0.10 Bonding in metals and metal-non-metal salts  (Page 7/9)

Why would size be a determining factor in the lattice energy? We should recall that the lattice energy follows Coulomb’s law. So, the closer the charges are to one another, the stronger is the interaction. Smaller ions can be closer together than larger ions. So the lattice energy is largest for the smallest ions.

Of course, Coulomb’s law also involves the number of the charges. In all of the compounds in [link] , the ions have a single +1 or -1 charge. We should look at compounds which contain doubly-charged ions. For common ions with +2 charges, we can look at the alkali earth metals In [link] , we can easily see that the lattice energies for salts of these ions are much larger than for the alkali metal ions. One final comparison would be a doubly-charged negative ion like O ^2- . Again, the lattice energies involving single positive charges with O ^2- are larger, and the lattice energy is even larger still when both ions are doubly charges, as in MgO.

We can conclude that compounds of metals and non-metals are typically formed by ionic bonding, and the strength of this bonding can be clearly understood using Coulomb’s law.

Observation 3: properties and bonding in solid carbon

In the first two observations of this study, we considered bonding in solids of two types, metals and salts. These are just two of the many types of solids, and not all solids are formed by either ionic bonding or metallic bonding. Far from it. We cannot look at every type of solid in this study, but it is worth considering one specific example which forms an interesting contrast to metals and salts. This example is diamond, one of several forms of pure solid carbon. (The other primary forms are graphene and the set of materials called fullerenes. We will postpone study of those materials for later.)

As always, we should begin with experimental observations to guide our understanding of diamond. What are its primary properties? It is a very hard solid, generally regarded as the hardest solid available in bulk. It is not malleable and would not be considered brittle like NaCl. It can be cleaved only with significant force. It has a very, very high melting point, over 3500 °C, and as an interesting note, it is the most thermally conducting material we know, meaning that it transfers heat better than any other substance. However, it does not conduct electricity.

What can we infer about the bonding in diamond from these properties? Since it is not brittle, we do not expect ionic bonding in diamond. This makes sense, since all of the atoms are carbon. But since it does not conduct electricity, we do not expect that the electrons are delocalized over the entire crystal, as they are in a metal. The bonding electrons must be more localized to individual nuclei. Since diamond is very hard and not malleable, the bonding must depend on the specific arrangement of the atoms, since, unlike a metal, it is very difficult to rearrange these atoms. Diamond won’t bend!