Diatomic halogens

Figure 9.5 Historical absorbance data for the construction of the Morse potential energy curves for the Halogen diatomic anions. The energies indicated by dotted lines were predicted in 1966 [12].

Infrared spectroscopy has broad appHcations for sensitive molecular speciation. Infrared frequencies depend on the masses of the atoms iavolved ia the various vibrational motions, and on the force constants and geometry of the bonds connecting them band shapes are determined by the rotational stmcture and hence by the molecular symmetry and moments of iaertia. The rovibrational spectmm of a gas thus provides direct molecular stmctural information, resulting ia very high specificity. The vibrational spectmm of any molecule is unique, except for those of optical isomers. Every molecule, except homonuclear diatomics such as O2, N2, and the halogens, has at least one vibrational absorption ia the iafrared. Several texts treat iafrared iastmmentation and techniques (22,36—38) and thek appHcations (39—42). 

In these examples tire entropy change does not vaty widely, and the value of the equilibrium constant is mainly determined by the heat of dissociation. It can be concluded, tlrerefore, that niuogen is one of the most stable diatomic molecules, and tlrat chlorine is tire most stable diatomic halogen molecule. 

The next step is to consider tire cross-sections of the absorption of radiation by the diatomic halogen molecules in order to decide if the relative effects result from the efficiency of the radiation photon-molecule interactions. These are reflected in the dissociation cross-sections of tlrese interactions. 

The halogens are volatile, diatomic elements whose colour increases steadily with increase in atomic number. Fluorine is a pale yellow gas which condenses to a canary yellow liquid, bp — 188.UC (intermediate between N2, bp —195.8°, and O2, bp — 183.0°C). Chlorine is a greenish-yellow gas, bp —34.0°, and bromine a dark-red mobile liquid, bp 59.5° interestingly the colour of both elements diminishes with decrease in temperature and at —195° CI2 is almost colourless and Br2 pale yellow. Iodine is a lustrous, black, crystalline solid, mp 113.6°, which sublimes readily and boils at 185.2°C. 

The molecular and bulk properties of the halogens, as distinct from their atomic and nuclear properties, were summarized in Table 17.4 and have to some extent already been briefly discussed. The high volatility and relatively low enthalpy of vaporization reflect the diatomic molecular structure of these elements. In the solid state the molecules align to give a layer lattice p2 has two modifications (a low-temperature, a-form and a higher-temperature, yS-form) neither of which resembles the orthorhombic layer lattice of the isostructural CI2, Br2 and I2. The layer lattice is illustrated below for I2 the I-I distance of 271.5 pm is appreciably longer than in gaseous I2 (266.6 pm) and the closest interatomic approach between the molecules is 350 pm within the layer and 427 pm between layers (cf the van der Waals radius of 215 pm). These values are... 

Figure 17.2 Schematic molecular orbital energy diagram for diatomic halogen molecules. (For F2 the order of the upper and 7T bonding MOs is inverted.).

All six possible diatomic compounds between F, Cl, Br and I are known. Indeed, ICl was first made (independently) by J. L. Gay Lussac and H. Davy in 1813-4 soon after the isolation of the parent halogens themselves, and its existence led J. von Liebig to miss the discovery of the new element bromine, which has similar properties (p. 794). The compounds vary considerably in thermal stability CIF is extremely robust ICl and IBr are moderately stable and can be obtained in very pure crystalline form at room temperature BrCl readily dissociates reversibly into its... 

Table 6-VI lists some properties of the halogens. In the elemental state, the halogens form stable diatomic molecules. This stability is indicated by the fact that it takes extremely high temperatures to disrupt halogen molecules to form the monatomic species. For example, it is known that the chlorine near the surface of the sun, at a temperature near 6000°C, is present as a gas consisting of single chlorine atoms. At more normal temperatures, chlorine atoms react with each other to form molecules ...

Apparently the diatomic molecules of the halogens already have achieved some of the stability characteristic of the inert gas electron arrangement. How is this possible How could one chlorine atom satisfy its need for one more electron (so it can reach the argon stability) by... 

FIGURE 2.20 Bond lengths fin picometers) of the diatomic halogen molecules. Notice how the bond lengths increase down the group as the atomic radii become larger. 

The halogens include fluorine, chlorine, bromine and iodine and all have been used in CVD reactions. They are reactive elements and exist as diatomic molecules, i.e., F2, CI2, etc. Their relevant properties are listed in Table 3.2. 

The first column of the periodic table, Group 1, contains elements that are soft, shiny solids. These alkali metals include lithium, sodium, potassium, mbidium, and cesium. At the other end of the table, fluorine, chlorine, bromine, iodine, and astatine appear in the next-to-last column. These are the halogens, or Group 17 elements. These four elements exist as diatomic molecules, so their formulas have the form X2 A sample of chlorine appears in Figure EV. Each alkali metal combines with any of the halogens in a 1 1 ratio to form a white crystalline solid. The general formula of these compounds s, AX, where A represents the alkali metal and X represents the halogen A X = N a C 1, LiBr, CsBr, KI, etc.). 

Table 922 lists average bond lengths for the most common chemical bonds. The table displays several trends. One trend is that bonds become longer as the radii of the atoms become larger. Compare the bond lengths of the diatomic halogens ... 

The halogens, the elements from Group 17 of the periodic table, provide an introduction to intermolecular forces. These elements exist as diatomic molecules F2, CI2, Bf2, and I2. The bonding patterns of the four halogens are identical. Each molecule contains two atoms held together by a single covalent bond that can be described by end-on overlap of valence p orbitals. 

In nature, the halogens exist as nonpolar diatomic molecules. London dispersion forces are the only forces of attraction acting between the molecules. These forces increase with increasing molecular size. 

F2 has the most positive standard reduction potential and therefore is the strongest of all common oxidizing agents. Oxidizing strengths of the diatomic halogen molecules decrease down Group 7A. 

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