pls, can someone tell me why the melting point of graphite is higher than that of diamond?

[bettilina]

pls can someone tell me why the melting point of graphite is higher than that of diamond? i need detailed a-level chemistry explanation

 

[matchwhiz]

pls can someone tell me why the melting point of graphite is higher than that of diamond? i need detailed a-level chemistry explanation

graphite has covalent bonding same as diamond,but it also have van der waal forces,which exist between the hexagonal parallel layers,thus more energy is needed in graphite due to the presence of van der waal forces,which are absent in diamond,thus graphite has a higher melting point than diamond.

 

[Doctor Nemo]

pls can someone tell me why the melting point of graphite is higher than that of diamond? i need detailed a-level chemistry explanation

 

[Doctor Nemo]

In graphite each carbon is linked to three others by covalent bonds. In diamond each carbon is linked to four others by covalent bonds. So it seems paradoxical that diamond would have a lower melting point. There are Van der Waals forces between the layers of graphite but these are relatively weak in comparison to a covalent bond. These weak forces allow the layers of graphite to slide past each other.

The layers of graphite contain delocalized electrons. There is one delocalized electron for each carbon atom. These delocalized electrons are shared by all the atoms in any one layer of graphite and make the bonds between the atoms stronger and the structure more stable in comparison to the bonds in diamond in the same way that delocalized electrons contribute to the greater stability of the benzene ring.

 

[aqeelhussain17]

In graphite each carbon is linked to three others by covalent bonds. In diamond each carbon is linked to four others by covalent bonds. So it seems paradoxical that diamond would have a lower melting point. There are Van der Waals forces between the layers of graphite but these are relatively weak in comparison to a covalent bond. These weak forces allow the layers of graphite to slide past each other.

The layers of graphite contain delocalized electrons. There is one delocalized electron for each carbon atom. These delocalized electrons are shared by all the atoms in any one layer of graphite and make the bonds between the atoms stronger and the structure more stable in comparison to the bonds in diamond in the same way that delocalized electrons contribute to the greater stability of the benzene ring.

In Graphite each C is bonded to three other C's? That something new for me, and I'm already done with Alevels. O: other than that everything seemed fine..

Sir sir, isnt like Diamond like the toughtest natural substance? Wow, all these days i thought diamond would have a higher melting point.. is all this info necesary for alevels? I never came across any question till now!

 

[bettilina]

graphite has covalent bonding same as diamond,but it also have van der waal forces,which exist between the hexagonal parallel layers,thus more energy is needed in graphite due to the presence of van der waal forces,which are absent in diamond,thus graphite has a higher melting point than diamond.

thanks, I already knw this but it isn't satisfactory to me, covalent bonds in diamond is very strong, is there any other explanation that you can give me?

 

[bettilina]

In graphite each carbon is linked to three others by covalent bonds. In diamond each carbon is linked to four others by covalent bonds. So it seems paradoxical that diamond would have a lower melting point. There are Van der Waals forces between the layers of graphite but these are relatively weak in comparison to a covalent bond. These weak forces allow the layers of graphite to slide past each other.

The layers of graphite contain delocalized electrons. There is one delocalized electron for each carbon atom. These delocalized electrons are shared by all the atoms in any one layer of graphite and make the bonds between the atoms stronger and the structure more stable in comparison to the bonds in diamond in the same way that delocalized electrons contribute to the greater stability of the benzene ring.

thanks sir, at least am getting a new light from the "benzene structure" direction

 

[Doctor Nemo]

thanks sir, at least am getting a new light from the "benzene structure" direction

Yes, each carbon in graphite is covalently bonded only to three other carbons in graphite. This leaves each carbon with an unpaired electron that can migrate in the 2-dimensional network. To cite the Wikipedia article on carbon,

“Because of the delocalization of one of the outer electrons of each atom to form a π-cloud, graphite conducts electricity, but only in the plane of each covalently bonded sheet. This results in a lower bulk electrical conductivity for carbon than for most metals. The delocalization also accounts for the energetic stability of graphite over diamond at room temperature.”

Assuming that this question is from a past paper, the CIE Chemistry syllabus requires knowledge of group 4 which includes carbon and its allotropes. I suppose you are expected to extrapolate from knowledge of stability of benzene due to its delocalized electrons to graphite. I don´t think it is an easy question.

 

[bettilina]

thanks so much

 

[MindStealth]

The melting point of graphite is slightly higher than the melting point of diamond because in graphite C-C bonds have partial double bond character and hence are stronger and difficult to break. In factor, carbon as graphite has the highest melting point among all the elements.
Most of the other answer...I don't know where they got that Idea from. No offence to anyone.