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202 Halogen Chemistry

the fluorine atom, its unusually high electronegativity The majority of commercially produced F 2 is used for
(Table III)—greater than that of any other element—and the manufacture of uranium hexafluoride (UF 6 , nuclear
the unavailability of low-energy d-orbitals in fluorine for power generation), sulfur hexafluoride (SF 6 , an important
chemical bonding. In spite of these differences, there are
closer similarities within the halogen family than within trifluoride (ClF 3 ), and the hexafluorides of tungsten and
any other family in the Periodic Table (with the exception rhenium (used for vapor deposition of the metal). Flu-
of the alkali metals, lithium through cesium). orine derivatives of hydrocarbons are used as refriger-
®
ants (Freon ), lubricants (Kel-F), and nonstick plastics
®
(Teflon ).
E. Properties of the Elemental States
Chlorine is produced industrially on a large scale in
Under ordinary conditions, the elemental halogens (ex- the “Chloralkai” process, electrolysis of aqueous sodium
cept astatine) exist as covalent, diatomic molecules, ac- chloride or natural brine solutions. The most commonly
quiring the electronic configuration of the next noble gas employed cell uses an asbestos diaphragm to separate the
by sharing one pair of electrons between two atoms. Their chlorine gas formed at the cathode from the sodium hy-
appearance ranges from a pale yellow gas (F 2 ), through a droxide (NaOH) that concentrates in the residual solution:
dark red liquid (Br 2 ), to an almost black, crystalline solid
2NaCl(aq) + 2H 2 O → Cl 2 (g) + H 2 (g) + 2NaOH(aq).
(I 2 ). Bromine is unique in being the only nonmetallic el-
(6)
ement that exists as a liquid under ordinary conditions.
Iodine crystals do not melt at atmospheric pressures, but Electrolysis of molten NaCl is also carried out on a large
sublime directly into the gas phase. scale, yielding sodium metal at the cathode rather than
The high volatilities and relatively low heats of vapor- hydrogen. Still other processes involve the electrolysis of
ization and fusion (Table IV) reflect the weak intermolec- fused magnesium chloride (MgCl 2 ) or the oxidation of
ular interactions expected for the covalent nature of these hydrogen chloride by oxygen or air in the presence of a
molecules. The stabilities of the X X bonds are indi- copper catalyst.
cated by their high dissociation energies (higher energies The major uses for elemental Cl 2 are the production
indicating stronger bonds). By comparison to the other of organic and inorganic compounds and bleaches. Some
halogens, the F F bond displays an anomalously low of the more important chlorinated compounds are used as
dissociation energy and large bond distance, attributable solvents, antifreeze, and plastics (e.g., polyvinyl chloride).
to electron–electron repulsion around the small fluorine Bromine is produced by oxidation of bromide ion using
atoms, or to the lack of a d-orbital component in the bond, Cl 2 gas (chlorine displacement):
or to both.
−
−
As indicated by their large, positive standard reduction 2Br (aq) + Cl 2 (g) → 2Cl (aq) + Br 2 (g). (7)
potentials, all of the halogen molecules have a strong ten-
Elemental bromine is blown out of the mixture as a vapor
dency to be reduced (i.e., to act as oxidizing agents). Flu-
using either steam or air and then condensed to form an
orine is the most powerful chemical oxidizing agent of all
impure product. Purification may involve reduction with
elements in the Periodic Table. This great reactivity is due
sulfur dioxide to reform the bromide ion,
primarily to the low energy of the F F bond. The general
tendency for halogen molecules to be reduced accounts Br 2 + SO 2 + 2H 2 O → 2HBr + H 2 SO 4 , (8)
for their natural occurrence as X or halide ions. The oxi-
−
followed by a second displacement [Eq. (7)] using
dizing power of fluorine is so great that this element does
chlorine gas. Alternatively, bromine can be added to
not normally exist in any oxidation state other than −1
sodium carbonate (Na 2 CO 3 ) solution to produce bromide
and 0; however, the remaining halogens can display ox-
and bromate:
idation states up to +7 if combined with a sufficiently
electronegative element such as oxygen. 3CO 2− + 3Br 2 → 5Br + BrO + 3CO 2 . (9)
−
−
3 3
On acidification, the bromine is regenerated:
F. Commercial Preparation and Use
+
−
−
5Br + BrO + 6H → 3Br 2 + 3H 2 O (10)
3
Because of its great reactivity, fluorine is prepared com-
mercially by electrolysis. The modern-day procedure is Until recently, almost all commercially prepared
similar to the original preparation by Moissan and utilizes bromine was converted to ethylene dibromide and used as
a 1:2 mixture of anhydrous KF and HF. The electrochem- a lead scavenger in gasoline. Other important applications
ical reaction produces hydrogen gas (H 2 ) at the cathode now include the production of pesticides, fire retardants,
and fluorine gas at the anode. The two gases will react drilling fluids, dyes, pharmaceuticals, and photographic
explosively and must be kept apart. chemicals.