To appreciate the chemistry of chlorine in comparison to that of fluorine it is necessary to appreciate the differences and trends between the elements. As may be seen in Table 1, chloride is significantly larger than fluorine. In addition while chlorine is an electronegative element its electronegativity is significantly less than that of fluorine, resulting in less polar bonding.

The X-Cl bond is an electron pair covalent bond with a highly polar nature. In this regard, chlorine is similar to fluorine. However, there are two key features with regard to chlorine’s bonding that differentiates it from fluorine.

The chloride ion (Cl^-) forms salts with ionic lattices (e.g., NaCl) but also forms a wide range of complexes, e.g., [Fe(H₂O)₅Cl]²⁺ and [RhCl₆]^3-. Chloride also acts as a bridging ligand in which one, two or three chlorides can bridge two metal centers (Figure 1).

Figure 1: Structures of common halide bridging moieties.

Chloride (and bromide) bridges are usually bent, whereas fluoride bridges can be either linear or bent. As an example, BeF₂ and BeCl₂ are isostructural, consisting of infinite chains with bent bridges (Figure 2). In contrast, transition metal pentahalides show different structures depending on the identity of the halide. This, TaCl₅ dimerizes with bent bridges (Figure 3a), while TaF₅ forms a cyclic tetramer with linear fluoride bridges (Figure 3b).

Figure 2: The solid state structure of BeX₂ (X = F or Cl).

Figure 3: The solid state structure of (a) TaCl₅ and (b) TaF₅.

Hydrogen chloride (HCl) is prepared by the reaction of concentrated sulfuric acid (H₂SO₄) with either NaCl or concentrated HCl solution.

Hydrogen chloride is a polar molecule with a dipole of 1.08 D. However, the lower polarity as compared to that of hydrogen fluoride (1.91 D) is consistent with the physical and chemical properties. Hydrogen chloride is a gas at room temperature (Mp = -114.25 °C, Bp = -85.09 °C), and its low boiling point is consistent with weak hydrogen bonding in the liquid state. While self-ionization, Equation 1, is very small, liquid HCl dissolves some inorganic compounds to give conducting solutions, Equation 2.

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Hydrogen chloride is soluble (and reacts) in water, Equation 3. The pK_a of the reaction (-7.0) is larger than observed for fluorine (3.2) and as such HCl is a stronger acid than HF.

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Chlorine forms a series of oxides (Table 2) in which the chlorine has the formal oxidation states +1, +4, +6, and +7. The physical properties of the oxides are summarized in Table 2. While, the oxides of chlorine are not very stable (in fact several are shock sensitive and are prone to explode) the conjugate oxyacids are stable.

Dichlorine monoxide (Cl₂O, Figure 4a) is a yellowish-red gas that is prepared by the reaction of chlorine with mercury oxide, Equation 4, or with a solution of chlorine in CCl₄.

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Figure 4: The structure of (a) Cl₂O, (b) ClO₂, (c) Cl₂O₄, and (d) Cl₂O₇.

When heated or subject to a spark, Cl₂O explodes to Cl₂ and O₂. Dichlorine monoxide reacts with water to form an orange-yellow solution of hypochlorous acid, Equation 5.

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Chlorine dioxide (ClO₂) is a yellowish gas at room temperature and is commonly used in industry as an oxidizing agent. The best synthesis of ClO₂ involves the reduction of potassium chlorate (KClO₃) by oxalic acid at 90 °C, since the CO₂ formed acts as a diluent for the highly explosive ClO₂. On an industrial scale ClO₂ is made by the exothermic reaction of sodium chlorate with SO₂ in sulfuric acid, Equation 6. The photolysis of ClO₂ yields a dark brown solid with the formula Cl₂O₃; however, its facile explosive decomposition precludes study.

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The structure of ClO₂ (Figure 4b) is equivalent to SO₂ with one extra electron, resulting in a paramagnetic unpaired electron species. Unusually, despite the unpaired electron configuration, ClO₂ shows no tendency to dimerize. This is unlike the analogous NO₂ molecule.

Dichlorine tetraoxide (Cl₂O₄) is commonly called chlorine perchlorate as a consequence of its structure (Figure 4c). Dichlorine hexaoxide (Cl₂O₆) is an unstable red oil that has the ionic structure in the solid state: [ClO₂]⁺[ClO₄]^-.

Dichlorine heptoxide (Cl₂O₇) is a relatively stable oil, that is prepared by the dehydration of perchloric acid at -10 °C, Equation 7, followed by vacuum distillation. The structure of Cl₂O₇ (Figure 4d) has been determined by gas phase electron diffraction.

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The reaction of Cl₂O₇ with alcohols and amines yields alkyl perchlorates (ROClO₃) and amine perchlorates (R₂NClO₃), respectively.

Given the isolobal relationship between the halogens it is not surprising that the mixed dihalogens can be prepared, e.g., ClF, ICl, and BrCl. Chlorine fluoride is a highly reactive gas (Bp = -100.1 °C) that is a powerful fluorinating agent, and is prepared by the oxidation of chlorine by chlorine trifluoride, Equation 8.

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The higher electronegativity of fluorine as compared to chlorine (Table 1), and the ability of chlorine to form more than one bond, means that higher fluorides of chlorine are also known, i.e., ClF₃ and ClF₅. Chlorine trifluoride (CF₃, Bp = 11.75 °C) is a useful fluorinating agent, that is prepared by the high temperature reaction of elemental chlorine and fluorine, is a useful fluorinating age. The gaseous pentafluoride (ClF₅, Bp = -31.1 °C) is prepared by the reaction of potassium chloride with fluorine, Equation 10.

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The structure of ClF₃ is T-shaped with two lone pairs on chlorine (Figure 5a), while that of ClF₅ is square pyramidal with a single lone pair on chlorine (Figure 5b).

Figure 5: The structures of (a) ClF₃ and (b) ClF₅.

In general the halogen fluorides are very reactive; explosive reactions occur with organic compounds. They are all powerful fluorinating agents when diluted with nitrogen, and the order of reactivity follows:

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Like most halogen fluorides, ClF, ClF₃ and ClF₅ all react with strong bases (e.g., alkali metal fluorides) to form anions, Equation 12 and Equation 13, and strong acids (e.g., AsF₅ and SbF₅) to form cations, Equation 14, Equation 15, and Equation 16.

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