Chapter 17: The Group 17 Elements

17.1 The Elements

Key Point: Except for fluorine and the highly radioactive astatine, the halogens exist with oxidation numbers ranging from −1 to +7; the small and highly electronegative fluorine atom is effective in oxidizing many elements to high oxidation states.

All the halogens have the valence electron configuration ns²np⁵. The features to note include their high ionization energies and their high electronegativities and electron affinities.

Property	F	Cl	Br	I	At
Covalent radius/pm	71	99	114	133	140
Ionic radius/pm	131	181	196	220	—
First ionization energy/kJ mol⁻¹	1681	1251	1139	1008	926
Pauling electronegativity	4.0	3.2	3.0	2.6	2.2
Electron affinity/kJ mol⁻¹	328	349	325	295	270
E°(X₂,X⁻)/V	+3.05	+1.36	+1.09	+0.54	—

Fluorine is a pale yellow gas that reacts with most inorganic and organic molecules and the noble gases Kr, Xe, and Rn. Chlorine is a green-yellow toxic gas. Bromine is the only liquid nonmetallic element at room temperature. Iodine is a purple-grey solid that sublimes to a violet vapor.

Anomaly of Fluorine:

F has a lower electron affinity than Cl, which seems at odds with its high electronegativity. This stems from larger electron-electron repulsion in the compact F atom compared to the larger Cl atom.

17.2 Simple Compounds

Key Point: All the halogens form hydrogen halides; HF is a liquid and HCl, HBr, and HI are gases. All the Group 17 elements form oxo compounds and oxoanions.

Hydrogen Halides

Property	HF	HCl	HBr	HI
Boiling point/°C	20	−85	−67	−35
Bond dissociation energy/kJ mol⁻¹	567	431	366	298
pK_a	3.45	c.−7	c.−9	c.−11

HF participates in extensive hydrogen bonding, giving it a wide liquid range and high relative permittivity. Although HF is a weak acid (pK_a = 3.45), it is one of the most toxic and corrosive substances known.

Halogen Oxides

Oxygen difluoride, OF₂, is the most stable oxide of F. Chlorine occurs with many different oxidation numbers in its oxides (+1, +3, +4, +6, +7). All chlorine oxides are endergonic and unstable.

Chlorine Oxoanions

Oxidation Number	Formula	Name	Shape
+1	ClO⁻	Hypochlorite	Linear
+3	ClO₂⁻	Chlorite	Angular
+5	ClO₃⁻	Chlorate	Pyramidal
+7	ClO₄⁻	Perchlorate	Tetrahedral

17.3 The Interhalogens

Key Point: All the halogens form compounds with other members of the group. Binary interhalogens have formulas XY, XY₃, XY₅, and XY₇, where the heavier, less electronegative halogen X is the central atom.

ClF₃

T-shaped (C_2v)

BrF₅

Square pyramidal (C_4v)

IF₇

Pentagonal bipyramidal (D_5h)

ICl

Linear

BrF₃

T-shaped

IF₅

Square pyramidal

The shapes of interhalogen molecules are largely in accord with the VSEPR model. XY₃ compounds have five valence electron pairs around X in a trigonal-bipyramidal arrangement, giving a bent T shape.

Example: BrF₃ Autoionization

2 BrF₃(l) ⇌ BrF₂⁺(sol) + BrF₄⁻(sol)

This Lewis acid-base behavior makes BrF₃ a useful solvent for ionic reactions under highly oxidizing conditions.

17.4 Occurrence, Recovery, and Uses

Key Point: Fluorine, chlorine, and bromine are prepared by electrochemical oxidation of halide salts; chlorine is used to oxidize Br⁻ and I⁻ to the corresponding dihalogen.

Fluorine Production

Elemental fluorine is produced by electrolysis of a 1:2 mixture of molten KF and HF. Aqueous electrolyte cannot be used because water is oxidized at a lower potential (+1.23 V).

Chloralkali Process

Most commercial chlorine is produced by electrolysis of aqueous NaCl solution:

Anode: 2 Cl⁻(aq) → Cl₂(g) + 2 e⁻
Cathode: 2 H₂O(l) + 2 e⁻ → 2 OH⁻(aq) + H₂(g)

Bromine and Iodine Recovery

Bromine is obtained by chlorine oxidation of Br⁻ in seawater:

Cl₂(g) + 2 X⁻(aq) → 2 Cl⁻(aq) + X₂(g) (X = Br or I)

One natural source of iodine is sodium iodate, NaIO₃. Which reducing agents would be thermodynamically feasible for iodine recovery?

17.5 Molecular Structure and Properties

Key Point: The F−F bond is weak relative to the Cl−Cl bond; bond strengths decrease down the group from chlorine.

Among the most striking physical properties of the halogens are their colors, ranging from almost colorless F₂ to purple I₂. The progression reflects the decrease in HOMO-LUMO gap down the group.

Bond Dissociation Enthalpies

Molecule	Bond Enthalpy (kJ mol⁻¹)
F₂	159
Cl₂	243
Br₂	193
I₂	151

Weak F−F Bond Explanation:

The low F−F bond enthalpy is consistent with low single-bond enthalpies of N−N and O−O. The bond is weakened by strong repulsions between nonbonding electrons in the small F₂ molecule.

17.6-17.8 Reactivity and Special Properties

Key Point: Fluorine is the most oxidizing halogen; the oxidizing power decreases down the group.

Pseudohalogens

Compounds with properties similar to halogens are called pseudohalogens. Examples include cyanogen (CN)₂ and thiocyanogen (NCS)₂.

Pseudohalide	Pseudohalogen	E°/V	Acid	pK_a
CN⁻	NCCN	+0.27	HCN	9.2
NCS⁻	NCSSCN	+0.77	HNCS	−1.9
N₃⁻	—	—	HN₃	4.92

Special Properties of Fluorine Compounds

Key Point: Fluorine substituents promote volatility, increase the strengths of Lewis and Brønsted acids, and stabilize high oxidation states.

High-oxidation-state fluorine compounds include IF₇, PtF₆, BiF₅, UF₆, and ReF₇. Fluorine also tends to disfavor low oxidation states—CuF is unstable but CuCl, CuBr, and CuI are stable.

17.10 The Interhalogens (Detail)

(a) Chemical Properties

Key Point: Fluorine-containing interhalogens are typically Lewis acids and strong oxidizing agents.

ClF₃ is an aggressive fluorinating agent:

2 Co₃O₄(s) + 6 ClF₃(g) → 6 CoF₃(s) + 3 Cl₂(g) + 4 O₂(g)

(b) Cationic Interhalogens

Under strongly oxidizing conditions, I₂ is oxidized to the blue paramagnetic diiodinium cation, I₂⁺. Higher polyhalogen cations include Br₅⁺, I₃⁺, and I₅⁺.

(c) Polyhalides

Polyiodides form when I₂ is added to I⁻ ions. The I₃⁻ ion is the most stable member of the series [(I₂)ₙI⁻]. Large cations stabilize polyiodides in the solid state.

Example: I₃⁻ Bonding

The Lewis structure of I₃⁻ has three equatorial lone pairs on the central I atom and two axial bonding pairs in a trigonal-bipyramidal arrangement, consistent with its linear structure.

17.11 Halogen Oxides

Key Point: The only fluorine oxides are OF₂ and O₂F₂; chlorine oxides are known for Cl oxidation numbers of +1, +4, +6, and +7; ClO₂ is the most commonly used halogen oxide.

Oxygen Difluoride (OF₂)

Prepared by passing fluorine through dilute aqueous hydroxide:

2 F₂(g) + 2 OH⁻(aq) → OF₂(g) + 2 F⁻(aq) + H₂O(l)

Chlorine Dioxide (ClO₂)

The only halogen oxide produced on a large scale. Used to bleach paper pulp and disinfect water:

2 ClO₃⁻(aq) + SO₂(g) → 2 ClO₂(g) + SO₄²⁻(aq)

Iodine Pentoxide (I₂O₅)

The most stable halogen oxide. Used to oxidize CO quantitatively to CO₂ in analytical chemistry.

17.12-17.15 Oxoacids and Redox Properties

Key Point: The halogen oxoanions are thermodynamically strong oxidizing agents; perchlorates of oxidizable cations are unstable.

Acidity of Chlorine Oxoacids

Acid	Formula	pK_a
Hypochlorous	HOCl	7.53 (weak)
Chlorous	HOClO	2.00
Chloric	HOClO₂	−1.2
Perchloric	HOClO₃	−10 (strong)

Thermodynamic Aspects

The Frost diagram shows that many intermediate oxidation states are susceptible to disproportionation. All oxidation states except the lowest (Cl⁻, Br⁻, I⁻) are strongly oxidizing.

Example: Chlorous Acid Disproportionation

2 HClO₂(aq) → ClO₃⁻(aq) + HClO(aq) + H⁺(aq) E°_cell = +0.52 V

Kinetic Trends

Key Point: Oxidation by halogen oxoanions is faster for lower oxidation states; rates and thermodynamics are both enhanced by acidic medium.

Rate order: ClO₄⁻ < ClO₃⁻ < ClO₂⁻ ≈ ClO⁻ ≈ Cl₂

Heavier halogens react faster: ClO₄⁻ < BrO₄⁻ < IO₄⁻

17.16 Fluorocarbons

Key Point: Fluorocarbon molecules and polymers are resistant to oxidation.

Synthesis Methods

Direct reaction with oxidizing metal fluoride:

RH(l) + 2 CoF₃(s) → RF(sol) + 2 CoF₂(s) + HF(sol)

Halogen exchange with catalyst:

CCl₄(l) + HF(l) → CCl₃F(l) + HCl(g)

PTFE (Polytetrafluoroethene)

Sold as Teflon®, PTFE is chemically inert, thermally stable (−196 to 260°C), and has low friction. The F atoms form a protective sheath around the carbon backbone.

n C₂F₄ → (−CF₂−CF₂−)ₙ

Applications of Fluorocarbons

Electrical tapes and coaxial cables (low conductivity)
Seals, piston rings, and bearings (mechanical properties)
Nonstick cookware coatings
Gore-Tex® fabric
HFCs as refrigerants (replacing ozone-depleting CFCs)

Summary: Key Reactions

Dihalogen Disproportionation

X₂(aq) + 2 OH⁻(aq) ⇌ XO⁻(aq) + X⁻(aq) + H₂O(l)

K = 7.5×10¹⁵ for Cl, 2×10⁸ for Br, 30 for I

Hypochlorite Disproportionation

3 ClO⁻(aq) ⇌ 2 Cl⁻(aq) + ClO₃⁻(aq) K = 1.5×10²⁷

Used for commercial production of chlorates.

Interhalogen Formation

ClF₃ + SbF₅ → [ClF₂]⁺[SbF₆]⁻

Strong Lewis acid SbF₅ abstracts F⁻ from interhalogen fluorides.