0.6 Molecular structure and physical properties  (Page 2/5)

To develop an understanding of bonding in these compounds, we focus on the halides of these elements. In , we compare physical properties of the chlorides of elements in Groups I and II to thechlorides of the elements of Groups IV, V, and VI, and we see enormous differences. All of the alkali halides and alkaline earthhalides are solids at room temperature and have melting points in the hundreds of degrees centigrade. The melting point of $\mathrm{Na}\mathrm{Cl}$ is 808°C, for example. By contrast, the melting points of the non-metal halides from Periods 2 and 3, such as $C{\mathrm{Cl}}_4$ , $P{\mathrm{Cl}}_3$ , and $S{\mathrm{Cl}}_2$ , are below 0°C, so that these materials are liquids at roomtemperature. Furthermore, all of these compounds have low boiling points, typically in the range of 50°C to 80°C.

Second, the non-metal halide liquids are electrical insulators, that is, they do not conduct an electricalcurrent. By contrast, when we melt an alkali halide or alkaline earth halide, the resulting liquid is an excellent electricalconductor. This indicates that these molten compounds consist of ions, whereas the non-metal halides do not.

We must conclude that the bonding of atoms in alkali halides and alkaline earth halides differs significantlyfrom bonding in non-metal halides. We need to extend our valence shell electron model to account for this bonding, and inparticular, we must account for the presence of ions in the molten metal halides. Consider the prototypical example of $\mathrm{Na}\mathrm{Cl}$ . We have already deduced that Cl atoms react so as to form acomplete octet of valence shell electrons. Such an octet could be achieved by covalently sharing the single valence shell electronfrom a sodium atom. However, such a covalent sharing is clearly inconsistent with the presence of ions in molten sodium chloride.Furthermore, this type of bond would predict that $\mathrm{Na}\mathrm{Cl}$ should have similar properties to other covalent chloride compounds, most of which are liquids at room temperature. Bycontrast, we might imagine that the chlorine atom completes its octet by taking the valence shell electron from a sodium atom,without covalent sharing. This would account for the presence of ${\mathrm{Na}}^{+}$ and ${\mathrm{Cl}}^{-}$ ions in molten sodium chloride.

In the absence of a covalent sharing of an electron pair, though, what accounts for the stability of sodiumchloride as a compound? It is relatively obvious that a negatively charged chloride ion will be attracted electrostatically to apositively charged sodium ion. We must also add to this model, however, the fact that individual molecules of $\mathrm{Na}\mathrm{Cl}$ are not generally observed at temperatures less than 1465°C, the boiling point of sodium chloride. Note that, if solid sodiumchloride consists of individual sodium ions in proximity to individual chloride ions, then each positive ion is not simplyattracted to a single specific negative ion but rather to all of the negative ions in its near vicinity. Hence, solid sodiumchloride cannot be viewed as individual $\mathrm{Na}\mathrm{Cl}$ molecules, but must be viewed rather as a lattice of positive sodium ions interacting with negative chloride ions. This type of“ionic” bonding, which derives from the electrostatic attraction of interlocking lattices of positive and negative ions,accounts for the very high melting and boiling points of the alkali halides.