Properties of Halogens

9 Key excerpts on "Properties of Halogens"

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  Inorganic Chemistry

  Butterworths Intermediate Chemistry

  • C. Chambers, A. K. Holliday(Authors)
  • 2016(Publication Date)
  • Butterworth-Heinemann

    (Publisher)

  272 11.1 Physical properties Table 11.1 and Table 11.2 (p. 274) give some of the physical properties of the common halogens. Figure 11.1 shows graphically some of the properties given in Table 11.1, together with enthalpies of atomization. All the elements exist as diatomic molecules X 2 . It can be seen that many properties change regularly with increasing atomic number, the changes being approximately linear in the case of the three elements chlorine, bromine and iodine, but a discontinuity almost always occurs for fluorine. This behaviour is typical for a group head element, which in addition tends to display properties not shown by other members of the group; a greater disparity in properties occurs between the first and second elements in a group than between any other two adjacent group elements. 11.1.1 Oxidation states The electronic configuration of each halogen is one electron less than that of a noble gas, and it is not surprising therefore, that a halogen atom can accept an electron to form X. Indeed, the reactions X(g) + e~ -X~(g)areall exothermic and the values (see Table 11.1), though small relative to the ionization energies, are all larger than the electron affinity of any other atom. Numerous ionic compounds with halogens are known but a noble gas configuration can also be achieved by the formation of a covalent bond, for example in halogen molecules, X 2 , and hydrogen halides, HX. When the fluorine atom acquires one additional electron the second quantum level is completed, and further gain of electrons is not energetically possible under normal circumstances, i.e. promotion to 3s requires too much energy. Thus fluorine is normally confined to a valency of 1 although in some solid fluorides bridge structures M—F—M are known in which fluorine acquires a covalency of 2. All the remaining halogens have unfilled d orbitals available and the covalency of the element can be increased.

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  The Chemistry of Chlorine, Bromine, Iodine and Astatine

  Pergamon Texts in Inorganic Chemistry, Volume 7

  • A. J. Downs, C. J. Adams(Authors)
  • 2016(Publication Date)
  • Pergamon

    (Publisher)

  In all cases, however, the electrical conductivity falls below the range characteristic of fused salts. Thus, all of the physical properties of the interhalogens are compatible with their formulation as molecular compounds. Accordingly, variations in the boiling and melting points and in the volatility of the compounds follow the pattern set by molecular polarizability (which is also a function of molecular geometry) and dipole moment (which spans the range from virtually zero for IF 7 to 2-18 D for IF 5 ). The properties of the interhalogens signify that the mononuclear unit ΧΥ Λ is commonly the only recognizable molecular aggregate in the solid, liquid or gaseous phases. A notable exception is provided, however, by iodine trichloride, crystals of which are composed of planar I2CI6 molecules (see below); the mononuclear unit ICI3 has not been characterized. Again, polymeric networks, which denote appreciable interaction between the IX units, have been established for crystalline iodine(I) chloride and bromide. According to the Trouton constants given in the table, chlorine(I) fluoride and iodine(I) chloride would appear also to be associated in the liquid state, but the experimental results may well be substantially in error. The non-ideality of the vapours of chlorine and bromine trifluorides have been attributed to the equilibrium 2XF 3 ^-X 2 F 6 for which equilibrium constants have been reported 838 ; thus, for the chlorine compound, K p (= P 2 CIFJPCUF*) i s 35' 4 atm at 24-2°C. When chlorine or bromine trifluoride is isolated at comparatively high concentrations in solid inert matrices maintained at 5-25°K, the infrared spectrum bears direct witness that dimers and other polymeric species exist as distinct, though weakly associated, aggregates 875 . Some degree of aggregation is also indicated by the comparatively large Trouton constants reported for the trifluorides.

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  Analytical Chemistry of Organic Halogen Compounds

  International Series in Analytical Chemistry

  • L. Mázor, R. Belcher, H. Freiser(Authors)
  • 2013(Publication Date)
  • Pergamon

    (Publisher)

  C H A P T E R I P R O P E R T I E S , P R E P A R A T I O N A N D R E A C T I O N S O F T H E H A L O G E N S A N D O F O R G A N I C H A L O G E N C O M P O U N D S I . T H E P R O P E R T I E S O F T H E H A L O G E N S 1. Physical properties There are five halogens, but only four of them have practical importance. At room temperature, fluorine and chlorine are gases, bromine is a liquid and iodine is a solid. Various gradual changes can be established in their physical properties. The most important physical characteristics are listed in Table I. The halogens are diatomic molecules in the solid, liquid and vapour states. The stability of the molecules decreases with atomic number; thus, for example, chlorine molecules are 5 % dissociated at 1500°C, while iodine exists exclusively in the monoatomic state at 1000°C. The halogens are the most electronegative elements. Their electronegativ-ity decreases with increasing atomic number, but even so, the electro-negativity of iodine is approximately the same as that of sulphur. There are seven electrons in the valence shell of halogens. Therefore, the maximum oxidation number of the halogens, except fluorine, is + 7 . Fluorine is the most electronegative element, and is found only in — 1 and 0 oxidation states. Each halogen occurs most frequently in the — 1 oxidation state, as they readily take up an electron to reach the next noble gas electron con-figuration. It is relatively difficult to perturb the valence electron shell of halogens, but it becomes easier with increasing atomic number. The value of the first ionization energy of fluorine is exceeded only by those of helium and neon. Halogens dissolve rather well in water; 1.46 g of chlorine is dissolved in 100 ml of water at 0°C; the respective values for bromine and iodine at 20°C are 3.58 g and 0.029 g, respectively. With fluorine, a violent reaction takes place. 2 * 19 20 ANALYTICAL CHEMISTRY OF ORGANIC HALOGEN COMPOUNDS 2.

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  Experimental Inorganic/Physical Chemistry

  An Investigative, Integrated Approach to Practical Project Work

  • M A Malati(Author)
  • 1999(Publication Date)
  • Woodhead Publishing

    (Publisher)

  8 The Halogens 8.1 INTRODUCTION The halogens (F, CI, Br, I and At) occupy group 17 of the Periodic Table. The atoms have 7 outer electrons and hence share a pair of electrons between two atoms forming a σ bond in the molecules Xj. The bonds become longer down the group. The bond dissociation enthalpy decreases from chlorine down the group. However, fluorine has a value close to that of iodine. As the Van der Waals forces increase with the size of the molecules, the boiling points of the elements increase. Thus fluorine and chlorine are gases whereas bromine is a liquid and iodine is a solid. These facts and the trend in the dissociation enthalpy explain the high reactivity of fluorine and the decrease in reactivity down the group. The halogen atom requires one electron to complete the stable configuration of the next noble gas. Hence the anions X' represent the most stable oxidation state of -I. The electron attachment energy becomes less negative from CI to I as expected from their atomic radii. However, F has a less negative value than CI. This behaviour and the anomalous dissociation enthalpy of F2 are ascribed to interelectronic repulsions in the compact 2p sub-shell of F. The halogens tend to form σ bonds with other non-metals as well as between themselves. E-X bonds (where Ε is a non-metal and X is a halogen) increase in length but decrease in strength from F to I. Because F cannot expand its valence shell beyond 8, it forms one E-F bond whereas the lower halogens have available d orbitals. Hence they can reach a maximum covalence of seven. Fluorine is also unique because of its high electronegativity (the highest in the Periodic Table) and its small radius. Thus it forms strong hydrogen bonds: F-H-E where E=F, Ο or N. The positive oxidation numbers of the halogens from +1 to +7 increase by two units.

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  The Chemistry of Nonaqueous Solvents VB

  Acid and Aprotic Solvents

  • J.J. Lagowski(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  28 The purity of the liquids can be checked with the aid of conductometry with a Kel-F cell. 29 For gases, microsublimation and various spectroscopies (infrared, Raman-laser, 19 F-NMR) are used. HI. PHYSICAL PROPERTIES In Downs and Adams' review article 4 a very exhaustive comparative table lists the physical properties of the interhalogens. For our purpose we choose to show the most useful data for solution chemists, namely, molecular weight, boiling point, melting point, vapor pressures; thermodynamic properties of the gas phase and of the liquid phase; mechanical properties such as density and viscosity; and electrical properties such as specific conductivity, dielectric constant, and dipole moment. The spectroscopic and structural properties are discussed in Section VI. A. Volatility The properties shown in Table II are necessary in choosing the correct solvent: the temperature range in which the interhalogens are liquids under reasonable pressure will indicate how they can be handled. From these data it can be seen that at room temperature all the chlorine fluorides are gases at atmospheric pressure. The monofluoride is the most volatile followed by C1F 5 and C1F 3 , which may be used as solvents, but obviously only at low temperatures or under high pressure. On the other hand IBr and IC1 3 are solids at 25°C. Iodine monochloride is the only one which has been used thoroughly for this purpose above 27.5°C.

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  Handbook of Water Analysis

  • Leo M.L. Nollet, Leen S. P. De Gelder, Leo M.L. Nollet, Leen S. P. De Gelder(Authors)
  • 2013(Publication Date)
  • CRC Press

    (Publisher)

  They can form a very large number of inorganic and organic compounds. Most of the halides can be classified into two categories. The fluorides and chlorides of many metallic elements, especially those belonging to the alkali metal and alkaline earth metal (except beryllium) families, are ionic compounds. Most of the halides of nonmetals such as sulfur and phospho-rus are covalent compounds. Fluorine, being the most electronegative of all reactive elements, occurs only with only 0 and − 1 oxidation numbers. Chlorine, bromine, and iodine can have oxidation numbers of − 1, 0, + 1, + 3, + 5, and + 7 in ions or molecules. Table 8.1 summarizes some of the most important char-acteristics of the stable halogens. In nature, because of their high reactivity, the halogens are always found combined with other ele-ments. Chlorine, bromine, and iodine occur most often as halides in sea water, in soil, and in minerals like halite (NaCl), sylvite (KCl), iodargite (AgI), and bromargyrite (AgBr). Chloride is a major anionic component of the biomass. Fluorine occurs in sparingly soluble mineral deposits such as fluorite and fluorspar (CaF 2 ), cryolite (Na 3 AlF 6 ), and fluorapatite (Ca 5 (PO 4 ) 3 F). The most easily oxidized halogen element, the iodine, is also found in iodates. The halogens are toxic materials. Their toxicity together with their reactivity decreases from fluorine to iodine. Except for astatine, the halogens are produced on an industrial scale and are used as reagents or oxidizing agents. Chlorine production is by far the largest. It is accomplished by electrochemical oxida-tion of aqueous sodium chloride solutions. Fluorine, however, cannot be obtained by electrochemical oxidation of aqueous solutions. Water decomposition would come at lower potential than the oxidation of fluoride. Even if an electrode with high overpotential could be found, the evolved fluorine would react immediately with the water content of the electrolysis cell.

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  Comparative Inorganic Chemistry

  • Bernard Moody(Author)
  • 2013(Publication Date)
  • Arnold

    (Publisher)

  23 Group VII: the halogens Fluorine, chlorine, bromine and iodine 9 The name halogen (Greek, F hals genon = sea-salt pro-2,7 ducing) was used by 17 Berzelius because he wished CI to indicate that chlorine, 2,8,7 bromine and iodine occurred 35 in the sea as salts. Sodium Br chloride is sea-salt. The !,8,18,7 halogens are a distinctive 53 family of diatomic non-I metallic elements showing a 2,8, 18,18,7 progressive gradation in 85 reactivity. Fluorine shows At many anomalies as the first 2,8, 18,32,18,7 member of the family, and indeed, it could be argued that fluorine is so distinctive that it is not a halogen, in the family sense, at all. The extreme non-metallic properties of fluorine are toned down with suc-cessive members of the family until some slight metallic characteristics appear with iodine. The radioactive element, astatine, is not described here. As the atomic number increases the elements darken in colour, becoming less volatile and less reactive. Fluorine is a pale yellow gas at room tem-perature, lighter in colour than chlorine which is a greenish-yellow gas. Bromine is a dark red liquid which gives off a dense red vapour, while iodine is a shiny grey crystalline solid, giving a violet vapour on heating. Fluorine, chlorine and the vapour of bromine are extremely irritant to nose and throat passages and are very poisonous. Chlorine was discovered by Scheele in 1774 but not recognized as an element. The gas was formed in a reaction between hydrochloric acid, or muriatic acid as it was known (Latin, muria = brine), and manganese(rv) oxide (pyrolusite). Because aqueous solutions of the new substance evolved oxygen in sunlight and because it was formed by oxidation of muriatic acid, Berthollet called the gas oxymuriatic acid in 1785. In a later series of experiments, Davy failed to show the presence of oxygen, concluding that the gas was an element. In 1810, he called it chlorine (Greek, chloros = greenish-yellow).

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  Chemistry and Analysis of Radionuclides

  Laboratory Techniques and Methodology

  • Jukka Lehto, Xiaolin Hou(Authors)
  • 2011(Publication Date)
  • Wiley-VCH

    (Publisher)

  , require radiochemical separations before measurement.

  Table 11.1 Important halogen radionuclides

  11.2 Physical and Chemical Properties of the Halogens

  Halogens are nonmetals and are located in group 17 of the periodic table with the electron structure [Ng]ns2 np5 . There are seven electrons in the outermost shell, that is they have the octet structure of the noble gases but with one electron missing. As a result, they readily form negative ions with a single charge and have the oxidation state of −I (F− , Cl− , Br− , I− , At− ). In the case of fluorine, −I is, in practice, the only oxidation state. The halogens are extremely electronegative, fluorine being the most electronegative element of all. Salts formed with these negative ions of single charge, the halides, are highly soluble, and solubility increases with the atomic number. In higher oxidation states, the halogens also form oxoacids. For example, in the oxidation state +V, iodine forms iodic acid, whose anion is iodate (IO3 − ), and in the oxidation state +VII it forms periodic acid, whose anion is periodate (IO4 − ). As seen from Figure 11.1 , iodide is the predominant chemical form. Iodate is present in oxidizing conditions in alkaline solutions. Formation of periodate requires oxidizing conditions more extreme than any found in natural systems. Chlorine most commonly appears in solution as the chloride ion, Cl− , and the perchlorate ion, ClO4 − , can only be found in very oxidizing conditions, again not found in natural systems (Figure 11.1 ). The halogens can also be in the form of gases (F2 , Cl2 , Br2 , I2 , At2

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  Concise Chemistry of the Elements

  • S C Siekierski, J Burgess(Authors)
  • 2002(Publication Date)
  • Woodhead Publishing

    (Publisher)

  with the explosive tendencies and ready hydrolysis, to NH 3 plus HCIO. of NC1 3 . Because of the small radius of the atom (and of the F-ion) HF is a relatiYely weak acid. in contrast to the rest of haloacids. This is just one of numerous examples of the consequences of particularly strong hydrogen-bonding to fluoride (as to o:·ygen). Both the unaYailability of d orbitals in fluorine and its high electronegativity prevent its forming oxoacids. 13.2 PROPERTIES OF Cl, Br, I and At Chlorine and bromine at sufficiently low temperatures, and iodine at room temperature, form layered solid phases by stacking together the singly bonded X 2 molecules. Solid chlorine and bromine are electric insulators whereas iodine is a two-dimensional semiconductor. with an energy gap (Eg) of about 1.3 eV (for silicon Eg = l. l eV). Under a pressure of 170 kbar iodine becomes a genuine metallic conductor. One can presume that if astatine were not a radioactive element with a short half-life, r 1 . 2 = 8. l h, it would be found to form a metallic phase. As explained in Section 6.5. the increase in metallic properties down the p block Groups or with pressure is the result of a decreasing energy gap between the filled bonding and nonbonding bands and the empty conduction (antibonding) band. It is interesting to note that on going from Cl 2 to 1 2 the ratio of the distances between the nearest and next nearest neighbour atoms increases from 0.6 to 0.78. This means Sec. 13.2) Properties of Cl, Br, I and At 123 that the bonds broaden and show a tendency to overlap. ,,·hich is a prerequisite for formation of a metallic phase. Under normal conditions chlorine is a greenish gas, bromine a dark red liquid and iodine a black solid which sublimes to gi,·e violet ·apour. The strong visible absorption spectra of Br 2 and 1 2 arise from transitions from the highest-energy filled n-antibonding orbitals to the lowest-energy empty CT-antibonding molecular orbital.

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