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17.6: Metal Halides - Structures and Energetics

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    Simple binary salts

    When considering the structures of simple binary salts only a small number of types are important. The simplest are for crystals with formula MX and include NaCl, CsCl and ZnS where the CN of the anion and cation are the same and may be 6, 8 or 4.

    Compounds with formulae of MX2 or M2X will have different CN's for the cation and anion. The most important of these are the fluorite and rutile structures.

    For the MX case, one reason why a substance might favour one form over another is due to the geometry of the packing of the spheres. What is needed is to be able to maximise the interactions of oppositely-charged ions while at the same time minimise the interactions of similarly-charged ions.
    The larger the difference in the sizes of the ions affects the packing of the larger ions around the smaller ions. It is relatively easy to calculate the radius ratio (r+/r-) and from this determine the limits for the various CN's.
    For the ZnS type structure with CN 4:4 the radius ratio is predicted to be in the range 0.22 - 0.41 while for the NaCl type structure with CN 6:6 it is predicted that the radius ratio will be within the range of 0.41 - 0.73. For CsCl structures with CN 8:8, the radius ratio is expected to be greater than 0.73. The Table below shows the values for a number of alkali halides. Experimentally it is found that the only examples under normal conditions of temperature and pressure to adopt the CsCl structure are CsCl, CsBr and CsI whereas several other salts were predicted to have this structure based on their radius ratios.

    Among the other factors that might influence the final structure is the interaction between ions in addition to the nearest neighbours. However the energies involved would not seem to be sufficient to alter the results as shown by the Madelung constants for NaCl and CsCl which are 1.74756 and 1.76267 respectively.

    Radius ratio values for alkali halides

    X- / M+ Li Na K Rb Cs
    F 0.44 0.70 0.98 0.92 0.81
    Cl 0.33 0.52 0.73 0.82 0.93
    Br 0.31 0.49 0.68 0.76 0.87
    I 0.28 0.44 0.62 0.69 0.78

    Transition Metal Halides

    The only stable pentahalide is VF5, which is readily hydrolysed and a strong Lewis acid. In the solid state it exists as an infinite chain polymer with cis-bridging fluorides but in the vapour phase it has a trigonal-bipyramid monomeric structure. M.P. 19.5° and B.P. 48.3°C.

    Tetrahalides are formed by Ti and V. The Ti tetrahalides are fairly unreactive in redox and halogenation chemistry, unlike the V compounds. VCl4 and VBr4 dissociate spontaneously under ambient conditions to VX3 and X2. They also tend to halogenate organic material.

    All trihalides of the elements from Ti to Cr are known. Mn(III) and Co(III) are too oxidising to coexist with any halide except F- under ambient conditions, whereas Ti(III) and V(III) are moderately strongly reducing. Chromium(III) is fairly stable toward both reduction or oxidation. There is a marked tendency toward decreasing ionic character on passing from left to right across the period and from the fluorides to the heavier halides. Ferric chloride and bromide show essentially covalent behaviour such as low MPs and solubility in donor organic solvents.

    Many trihalides can be prepared by direct combination of the elements. In those cases where direct combination gives a higher oxidation state, trihalides can be produced by either thermal dissociation, disproportionation of the higher halide or by reduction- for example TiCl3 can be prepared by reduction of TiCl4 with H2 at high temperatures.

    All these trihalides adopt structures in which the metal is six-coordinate, either octahedral or distorted octahedral. Many of the lattices are complicated, but can be represented as CrCl3 or BiI3 types. In some cases it becomes even more complicated since some salts exist in more than one form. For example, at low temperature (below 240K) CrCl3 exists in the rhombohedral form mentioned above, but at room temperature it is monoclinic.

    All the first row transition metal elements form dihalides with all the halogens, with the exceptions of TiF2 and CuI2. The instability of TiF2 is probably due to easy disproportionation to Ti and TiF3 whereas the oxidising power of Cu2+ (reducing power of I-) explains the lack of the copper salt. Anhydrous dihalides can generally be synthesised by reaction of the pure metal with hydrogen halide or, for labile metal ions, by dehydration of hydrated salts with a covalent halogen compound e.g. SOCl2.

    The difluorides commonly have rutile structures, for example MnF2, the dichlorides CdCl2 structures and the diiodides CdI2 structures. Dibromides have either CdCl 2 or CdI2 structures or both. Dihalides are all ionic and typically dissolve in water to give aquo complexes or mixed aquo-halo-complexes. The solutions of Ti(II), V(II) and Cr(II) are very strongly reducing. They react extremely rapidly with O 2, and Ti(II) even rapidly reduces water to liberate hydrogen. Solutions of Fe(II) undergo slow oxidation in air, but in acid or neutral solution Mn(II), Co(II), Ni(II), and Cu(II) are quite stable to oxygen.

    CuI adopts the Zinc Blende cubic close packed structure.

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    17.6: Metal Halides - Structures and Energetics is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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