Why is diamond a solid, whereas carbon dioxide, which is a heavier molecule, is a gas?

1 Answer
Dec 4, 2016

Answer:

Because diamond is a non-molecular material, with no molecular boundaries.

Explanation:

Carbon dioxide is a molecular species, and occurs as discrete molecules of #O=C=O#. The forces of attraction between its constituent molecules are negligible.

On the other hand, diamond is a non-molecular solid. Its particles are held together by (very!) strong #C-C# (covalent!) bonds that persist across the entire lattice. As a result, the melting points/boiling points of diamond are so high as to be almost unmeasurable. This reflects the stability and strength of the lattice.

The important criterion is NOT whether the bonds are covalent. And in fact covalent bonds are strong; the covalent bonds in carbon dioxide, and carbon monoxide ARE IN FACT STRONGER than the #C-C# bonds in graphite or diamond. Material properties are determined by whether the material is #"molecular"#. Silicon dioxide has weaker #Si-O# bonds than the #Si-H# bonds of silane. Nevertheless, #SiO_2# is a very high melting solid in that its bonds are non-molecular and persist thruout the crystalline lattice. #SiH_4# is molecular, and is a gas.

What does this argument suggest with respect to the properties of ionic solids, such as #NaCl# and #CaF_2#? These are ionic solids, and what could we say, #"a priori"#, with regard to their molecularity, and physical properties?

See here for a consideration of metallic bonding.