Chemistry » Metals, Metalloids, and Nonmetals » Occurrence, Preparation, and Properties of Halogens

# Occurrence and Preparation of Halogens

The elements in group 17 are the halogens. These are the elements fluorine, chlorine, bromine, iodine, and astatine. These elements are too reactive to occur freely in nature, but their compounds are widely distributed. Chlorides are the most abundant; although fluorides, bromides, and iodides are less common, they are reasonably available.

In this section, we will examine the occurrence, preparation, and properties of halogens. Next, we will examine halogen compounds with the representative metals followed by an examination of the interhalogens. This section will conclude with some applications of halogens.

## Occurrence and Preparation

All of the halogens occur in seawater as halide ions. The concentration of the chloride ion is 0.54 M; that of the other halides is less than 10–4M. Fluoride also occurs in minerals such as CaF2, Ca(PO4)3F, and Na3AlF6. Chloride also occurs in the Great Salt Lake and the Dead Sea, and in extensive salt beds that contain NaCl, KCl, or MgCl2.

Part of the chlorine in your body is present as hydrochloric acid, which is a component of stomach acid. Bromine compounds occur in the Dead Sea and underground brines. Iodine compounds are found in small quantities in Chile saltpeter, underground brines, and sea kelp. Iodine is essential to the function of the thyroid gland.

The best sources of halogens (except iodine) are halide salts. It is possible to oxidize the halide ions to free diatomic halogen molecules by various methods, depending on the ease of oxidation of the halide ion. Fluoride is the most difficult to oxidize, whereas iodide is the easiest.

The major method for preparing fluorine is electrolytic oxidation. The most common electrolysis procedure is to use a molten mixture of potassium hydrogen fluoride, KHF2, and anhydrous hydrogen fluoride. Electrolysis causes HF to decompose, forming fluorine gas at the anode and hydrogen at the cathode. It is necessary to keep the two gases separated to prevent their explosive recombination to reform hydrogen fluoride.

Most commercial chlorine comes from the electrolysis of the chloride ion in aqueous solutions of sodium chloride; this is the chlor-alkali process discussed previously. Chlorine is also a product of the electrolytic production of metals such as sodium, calcium, and magnesium from their fused chlorides. It is also possible to prepare chlorine by the chemical oxidation of the chloride ion in acid solution with strong oxidizing agents such as manganese dioxide (MnO2) or sodium dichromate (Na2Cr2O7). The reaction with manganese dioxide is:

$${\text{MnO}}_{2}(s)+2{\text{Cl}}^{\text{−}}(aq)+4{\text{H}}_{3}{\text{O}}^{\text{+}}(aq)\;⟶\;{\text{Mn}}^{2+}(aq)+{\text{Cl}}_{2}(g)+6{\text{H}}_{2}\text{O}(l)$$

The commercial preparation of bromine involves the oxidation of bromide ion by chlorine:

$$2{\text{Br}}^{\text{−}}(aq)+{\text{Cl}}_{2}(g)\;⟶\;{\text{Br}}_{2}(l)+2{\text{Cl}}^{\text{−}}(aq)$$

Chlorine is a stronger oxidizing agent than bromine. This method is important for the production of essentially all domestic bromine.

Some iodine comes from the oxidation of iodine chloride, ICl, or iodic acid, HlO3. The commercial preparation of iodine utilizes the reduction of sodium iodate, NaIO3, an impurity in deposits of Chile saltpeter, with sodium hydrogen sulfite:

$$2{\text{IO}}_{3}{}^{\text{−}}(aq)+5{\text{HSO}}_{3}{}^{\text{−}}(aq)\;⟶\;3{\text{HSO}}_{4}{}^{\text{−}}(aq)+2{\text{SO}}_{4}{}^{2-}(aq)+{\text{H}}_{2}\text{O}(l)+{\text{I}}_{2}(s)$$