Silicon dioxide, or SiO2, is a large covalent oxide. It has a very high melting point. This is due to the large number of covalent bonds that are present in the molecule. Because of this, the molecule requires a lot of heat energy to melt.
To understand this phenomenon, let’s take a look at the chemistry of the molecule. Each silicon atom is surrounded by four oxygen atoms. The oxygen atoms are electronegative. These bonds give the compound unique properties. They are called “bridge” bonds. There are hundreds of these bonds in the molecule.
As you might expect, the more bonds there are, the higher the melting point. But there are other factors that play a role in this process. Specifically, the shape and size of the molecule determine how much van der Waals dispersion is present.
The strength of the bonds between the oxygen atoms and the silicon atom affects the melting and boiling points. If the bonds are strong, more energy is needed to break them. On the other hand, if the bonds are weak, the molecules are easy to break together.
In order to dissociate, the silicates will undergo a series of phase transitions. These include the a-PbO2-type, the CaCl2-type, and the stishovite phase. All of these phase transitions require a large amount of energy to break free from the lattice. And this is why the melting point of SiO2 is so high.
Currently, there is no definitive model of the solid-liquid boundary for SiO2 at pressures above 500 GPa. However, simulations have reported melting curves up to 3400 GPa25 using a two-phase method.