Introduction group VIIA

Group VIIA includes following elements - the halogens: Fluorine - [F], Chlorine - [Cl], Bromine - [Br], Iodine - [I], Astatine - [At].

Compounds of the halogens have been known from earliest times and the elements have played a particularly important role during the past two hundred years in the development of both experimental and theoretical chemistry. The name “halogen” was introduced by J. S. C. Schweigger in 1811 to describe the property of chlorine, at that time unique among the elements, of combining directly with metals to give salts. The name has since been extended to cover all five members of Group VIIa (Group 17) of the periodic table.


Fluorine derives its name from the early use of fluorspar (CaF2) as a flux (Latinfluor, flowing). The name was suggested to Sir Humphry Davy by A.M. Ampkre in 1812. The corrosive nature of hydrofluoric acid and the curious property that fluorspar has of emitting light when heated (“fluorescence”) were discovered in the seventeenth century. However, all attempts to isolate the element either by chemical reactions or by electrolysis were foiled by the extreme reactivity of free fluorine. Success was finally achieved on 26 June 1886 by H. Moissan who electrolysed a cooled solution of KHF2 in anhydrous liquid HF, using Pt/Ir electrodes sealed into a platinum U-tube sealed with fluorspar caps: the gas evolved immediately caused crystalline silicon to burst into flames.

Fluorine technology and the applications of fluorine-containing compounds have developed dramatically during the twentieth century. Noteworthy events are the development of inert fluorinated oils, greases and polymers: Freon gases such as CCl2F2 (1928) were specifically developed for refrigeration engineering; others were used as propellants in pressurized dispensers and aerosols; and the non-stick plastic polytetrafluoroethylene (PTFE or Teflon) was made in 1938. Inorganic fluorides, especially for the aluminium industry have been increasingly exploited from about 1900, and from 1940 UF6 has been used in gaseous diffusion plants for the separation of uranium isotopes for nuclear reactor technology. The great oxidizing strength of F2 and many of its compounds with N and O have attracted the attention of rocket engineers and there have been growing largescale industrial applications of anhydrous HF.

The aggressive nature of HF fumes and solutions has been known since Schwanhard of Nurnberg used them for the decorative etching of glass. Hydrofluoric acid inflicts excruciatingly painful skin bums and any compound that might hydrolyse to form HF should be treated with great caution. Maximum allowable concentration for continuous exposure to HF gas is 2-3 ppm (cf. HCN 10 ppm). The free element itself is even more toxic, maximum allowable concentration for a daily 8-h exposure being 0.1 ppm. Low concentrations of fluoride ion in drinking water have been known to provide excellent protection against dental caries since the classical work of H. T. Dean and his colleagues in the early 1930s; as there are no deleterious effects, even over many years, providing the total fluoride ion concentration is kept at or below 1 ppm, fluoridation has been a recommended and adopted procedure in several countries for many years. However, at 2-3ppm a brown mottling of teeth can occur and at 50 ppm harmful toxic effects are noted. Ingestion of 150mg of NaF can cause nausea, vomiting, diarrhoea and acute abdominal pains though complete recovery is rapid following intravenous or intramuscular injection of calcium ions.


Chlorine was the first of the halogens to be isolated and common salt (NaCl) has been known from earliest times. Its efficacy in human diet was well recognized in classical antiquity and there are numerous references to its importance in the Bible. On occasion salt was used as part payment for the services of Roman generals and military tribunes (salary) and, indeed, it is an essential ingredient in mammalian diets. The alchemical use of aqua regia (HCl/HNO3) to dissolve gold is also well documented from the thirteenth century onwards. Concentrated hydrochloric acid was prepared by J. L. Glauber in 1648 by heating hydrated ZnCl2 and sand in a retort and the pure gas, free of water, was collected over mercury by J. Priestley in 1772. This was closely followed by the isolation of gaseous chlorine by C. W. Scheele in 1774: he obtained the gas by oxidizing nascent HCl with MnO2 in a reaction which would now formally be written as:

4NaCl+ 2H2SO4 + MnO2 → 2Na2SO4 + MnCl2 + 2H2O + Cl2

The bleaching power of Cl2 was discovered by Scheele in his early work (1774) and was put to technical use by Berthollet in 1785. This was a major advance on the previous time-consuming, labour-intensive, weather-dependent method of solar bleaching, and numerous patents followed. Indeed, the use of chlorine as a bleach remains one of its principal industrial applications (bleaching powder, elemental chlorine, hypochlorite solutions, chlorine dioxide, chloramines, etc.). Another all-pervading use of chlorine, as a disinfectant and germicide, also dates from this period (1801), and the chlorination of domestic water supplies is now almost universal in developed countries. Again, as with fluoride, higher concentrations are toxic to humans: the gas is detectable by smell at 3 ppm, causes throat irritation at 15 ppm, coughing at 30 ppm, and rapid death at 1000 ppm. Prolonged exposure to concentrations above 1 ppm should be avoided.

Sodium chloride, by far the most abundant compound of chlorine, occurs in extensive evaporite deposits, saline lakes and brines, and in the ocean. It has played a dominant role in the chemical industry since its inception in the late eighteenth century. The now defunct Leblanc process for obtaining NaOH from NaCl signalled the beginnings of large-scale chemical manufacture, and NaCl remains virtually the sole source of chlorine and hydrochloric acid for the vast present-day chlorine-chemicals industry. This embraces not only the large-scale production and distribution of Cl2 and HCl, but also the manufacture of chlorinated methanes and ethanes, vinyl chloride, aluminium trichloride catalysts and the chlorides of Mg, Ti, Zr, Hf, etc., for production of the metals. About 15 000 chlorinated compounds are currently used to varying degrees in commerce. Of these, the environmental and health hazards posed by certain polychlorinated hydrocarbons is now well established, though not all such compounds are dangerous: focused selective restrictions rather than a blanket banning of all organochlorine compounds is advocated.


The magnificent purple pigment referred to in the Bible and known to the Romans as Tyrian purple after the Phoenician port of Tyre (Lebanon), was shown by P. Friedlander in 1909 to be 6,6’-dibromoindigo. This precious dye was extracted in the early days from the small purple snail Murex brunduris, as many as 12000 snails being required to prepare 1.5g of dye. The element itself was isolated by A.J. Balard in 1826 from the mother liquors remaining after the crystallization of sodium chloride and sulfate from the waters of the Montpellier salt marshes; the liquor is rich in MgBr2, and the young Balard, then 23 y of age, noticed the deep yellow coloration that developed on addition of chlorine water. Extraction with ether and KOH, followed by treatment of the resulting KBr with H2SO4/MnO2, yielded the element as a red liquid. Astonishingly rapid progress was possible in establishing the chemistry of bromine and in recognizing its elemental nature because of its similarity to chlorine and iodine (which had been isolated 15 y earlier). Indeed, J. von Liebig had missed discovering the element several years previously by misidentifying a sample of it as iodine monochloride. Balard had proposed the name muride, but this was not accepted by the French Academy, and the element was named bromine because of its unpleasant, penetrating odour. It is perhaps ironic that the name fluorine had already been preempted for the element in CaF2 and HF since bromine, as the only non-metallic element that is liquid at room temperature, would preeminently have deserved the name.

The first mineral found to contain bromine (bromyrite, AgBr) was discovered in Mexico in 1841, and industrial production of bromides followed the discovery of the giant Stassfurt potash deposits in 1858. The major use at that time was in photography and medicine: AgBr had been introduced as the light-sensitive agent in photography about 1840, and the use of KBr as a sedative and anti-convulsant in the treatment of epilepsy was begun in 1857. Other major uses of bromine-containing compounds include their application as flame retardants and as phase-transfer catalysts.


The lustrous, purple-black metallic sheen of resublimed crystalline iodine was first observed by the industrial chemist B. Courtois in 1811, and the name, proposed by J.L. Gay Lussac in 1813, reflects this most characteristic property (Greek, violet-coloured). Courtois obtained the element by treating the ash of seaweed (which had been calcined to extract saltpetre and potash) with concentrated sulfuric acid. Extracts of the brown kelps and seaweeds Fucus and Laminaria had long been known to be effective for the treatment of goitre and it was not long before J. F. Coindet and others introduced pure KI as a remedy in 1819.

The first iodine-containing mineral (AgI) was discovered in Mexico in 1825 but the discovery of iodate as an impurity in Chilean saltpetre in 1840 proved to be more significant industrially. The Chilean nitrate deposits provided the largest proportion of the world’s iodine until overtaken in the late 1960s by Japanese production from natural brines.

In addition to its uses in photography and medicine, iodine and its compounds have been much exploited in volumetric analysis (iodometry and iodimetry). Organoiodine compounds have also played a notable part in the development of synthetic organic chemistry, being the first compounds used in A. W. von Hofmann’s alkylation of amines (1850), A. W. Williamson’s synthesis of ethers (1851), A. Wurtz’s coupling reactions (1855) and V. Grignard’s reagents (1900).


From its position in the periodic table, all isotopes of element 85 would be expected to be radioactive. Those isotopes that occur in the natural radioactive series all have halflives of less than 1 min and thus occur in negligible amounts in nature. Astatine (Greek, unstable) was first made and characterized by D. R. Corson, K. R. Mackenzie and E. Segrk in 1940: they synthesized the isotope 211At (t1/2 ~7.21h) by bombarding 209Bi with α-particles in a large cyclotron:

20983Bi + 42He → 21185At + 221n

In all, some 27 isotopes from 194At to 220At have now been prepared by various routes but all are short-lived. The only ones besides 211At having half-lives longer than 1 h are 207At (1.80 h), 208At (1.63 h), 209At (5.41 h), and 210At (8.1 h): this means that weighable amounts of astatine or its compounds cannot be isolated, and nothing is known of the bulk physical properties of the element. For example, the least-unstable isotope (210At) has a specific activity corresponding to 2 curies per μg, i.e. 7 x 1010 disintegrations per second per μg. The largest preparations of astatine to date have involved about 0.05 μg and our knowledge of the chemistry of this element comes from extremely elegant tracer experiments, typically in the concentration range 10-10...10-15 mole.