Simple molecules contain only a few atoms held together by covalent bonds . An example is carbon dioxide (CO 2), the molecules of which contain one atom of carbon bonded with two atoms of oxygen.

A giant covalent structure is a three-dimensional structure of atoms that are joined by covalent bonds. Allotropes are different forms of the same element, in the same state. Graphite ,graphene and diamond are allotropes of the same element (carbon) in the same state (solid). Carbon can form up to four covalent bonds.

Diamond and graphite forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures (lattices) of atoms. All the atoms in these structures are linked to other atoms by strong covalent bonds and so they have very high melting points.

Silicon is metalloid and it is also considered as nonmetallic and chlorine is also a nonmetal and the bond between metals and non metal is ionic as well as bond between nonmetals and nonmetals are covalent. So SiCl4 is covalent.

Like graphene, nanotubes are strong, and they conduct electricity because they have delocalised electrons. Buckyballs are spheres or squashed spheres of carbon atoms. They are made up of large molecules but do not have a giant covalent structure.

Graphene has a very high melting point and is very strong because of its large regular arrangement of carbon atoms joined by covalent bonds .

Diamond is a giant covalent structure in which: each carbon atom is joined to four other carbon atoms by strong covalent bonds. the carbon atoms form a regular tetrahedral network structure. there are no free electrons.

Sulphur is a non-metal because it is consistent with three physical properties listed for non-metals. It is a poor conductor of heat and electricity, because electrons are not free to move.

Its molecules are made up of 60 carbon atoms joined together by strong covalent bonds. Molecules of C 60 are spherical. There are weak intermolecular forces between molecules of buckminsterfullerene. These need little energy to overcome, so buckminsterfullerene is slippery and has a low melting point.

2 Answers. To reiterate Ivan’s comment fullerene is a bad conductor because that’s what the measured properties produce as a result. The mechanism that makes it a bad conductor is that it has shorter range continuity than graphite. In graphite the carbon is made of sheets that can be as long as the sample.

All fullerenes have a structure similar to graphite in that each carbon atom forms single covalent bonds with three neighbouring carbon atoms. This forms a hexagonal plane structure, which can be bent to form different fullerenes. These can move throughout the fullerene, allowing conduction of electricity.

Diamond has carbon-carbon bonds that are covalent in nature, meaning they’re shared specifically between two carbon atoms. They are not free to move about the solid. This is why diamond is an electrical insulator.

For all isolated-pentagon isomers of fullerenes but C60(Ih), the minimum cation LE in a molecule (min cation LE) is smaller than the graphite value. Thus, fullerenes are predicted to be much more reactive than graphite. In fact, they are very susceptible to nucleophilic attack.

The remaining electron at each carbon is delocalised in molecular orbitals, which in turn give an aromatic character to the molecule. The pure form of fullerene acts as an insulator but can be converted into semiconductor or superconductor by doping under suitable conditions.

In fact C60 is diamagnetic (no unpaired electrons) and does not conduct electricity.

When C60 is compressed at 3 GPa and heated to 700°C it produces a form of carbon that is semimetallic and has a hardness that is approximately two-thirds that of diamond.

A single C60 fullerene also has delocalised electrons, but electrons cannot jump between individual fullerenes. So from a large scale, a that block of C60’s is not conductive. Graphene only conducts electricity very well down the sheet. It can’t conduct electricity well from the top to bottom direction.