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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 Na Cl is 808°C, for example. By contrast, the melting points of the non-metal halides from Periods 2 and 3, such as C Cl 4 , P Cl 3 , and S 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.

Melting points and boiling points of chloride compounds
Melting Point (°C) Boiling Point (°C)
Li Cl 610 1382
Be Cl 2 405 488
C Cl 4 -23 77
N Cl 3 -40 71
O Cl 2 -20 4
F Cl -154 -101
Na Cl 808 1465
Mg Cl 2 714 1418
Si Cl 4 -68 57
P Cl 3 -91 74
S Cl 2 -122 59
Cl 2 -102 -35
K Cl 772 1407
Ca Cl 2 772 >1600

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 Na 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 Na 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 Na + and 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 Na 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 Na 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.

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Source:  OpenStax, General chemistry i. OpenStax CNX. Jul 18, 2007 Download for free at http://cnx.org/content/col10263/1.3
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