Questions on structure and bonding

How come in small covalent structures, the melting and boiling points are low due to the weak intermolecular forces, but in giant covalent structures, they melting and boiling points are high, despite the intermolecular forces still being weak?

Small Covalent molecules when in liquid form are joined by the IMF.... making them liquid, not gaseous. When they reach their boiling point, these forces break and the molecules separate and rise into a gas. In giant covalent structures, the molecules are very large.... much larger than CO2 meaning they have more covalent bonds.
For example, each carbon in diamond has 4 bonds.... that means the carbon you look at bonded to that carbon will also have 4 bonds, and so on.
Whereas CO2 for example, only has 2 covalent bonds. It's carbon is bonded to two oxygens and those oxygens aren't bonded to another carbon.... unlike in diamond, but they hold an IMF with another CO2 molecule not a covalent bond, when non-gaseous.

Basically, giant structures have more covalent bonds than small covalent.

CO₂ doesn't have 2 covalent bonds - it has 2 $\pi$ bonds and 2 $\sigma$ bonds to give a total of 4 covalent bonds.

Yes apologies, 4 covalent bonds but the overall structure still has less covalent bonds.

Covalent bonds and intermolecular forces aren't the same thing. Giant covalent structures only really contain covalent bonding - graphite is an exception where its layers of graphene have intermolecular forces between them, but the amount of covalent bonding compared to intermolecular forces in graphite is very large. Simple covalent structures contain a lot more intermolecular forces.

Let's put it this way - giant covalent structures contain an extremely large repetition of atoms joined by covalent bonding. For example, diamond is pretty much entirely covalently bonded. To melt diamond, there are no bonds you can break other than each covalent bond, and this takes a lot of energy. This means diamond has a large melting point, and a large boiling point too.

However, simple (small) covalent molecules are different. An example of a carbon-based, simple covalent molecule is methane. The only covalent bonding which exists in a sample of methane is in each individual molecule (so, between each carbon and hydrogen atom), but the thing that keeps each methane molecule together are intermolecular forces, not covalent bonding. The intermolecular forces in methane are called London forces, but there are other types. However, in general, intermolecular forces are nowhere near as strong as covalent bonding, hence to melt a sample of methane, all you have to do is overcome the intermolecular forces. Compared to a giant covalent structure like diamond, this doesn't require as much energy, and so the melting point is lower, along with a lower boiling point.

I hope this helps.