If Sp^2 bonds are what makes nanotubes so strong, how come graphite with sp^2 bonds not stronger than diamonds

Question

So I understand that diamonds are stronger than graphite, but diamonds have sp^3 bonds and graphite has sp^2 bonds. And supposedly, nanotubes with sp^2 bonds similiar to graphite are pretty darn strong too. How come graphite isn't stronger than diamonds with its sp^2 bonds?

Answer 1

Graphite is made up of "sheets" of carbon atoms bonded to one another by sp2 bonds. The bonding *within* these sheets is actually *stronger* than the sp3 bonds in diamonds. As evidence of this, consider that the melting point of graphite is actually higher than that of diamond (see first source). To melt a substance, one has to disrupt the bonds between the individual molecules (or atoms, in this case) of the material.

Graphite is considered a "soft" material because the individual sheets of carbon atoms are very weakly bonded to one another by what are called "dispersion forces" (specifically, van der Waals forces). Random fluctuations in the electron density in the pi bonds of the graphite sheets cause the sheets to have small, temporary dipole moments. Interaction of the dipole moments of adjacent sheets causes them to stick together, but only very weakly. See second source.

Answer 2

sp3 bonds of diamond are single bonds extended out in four directions from every carbon atom. If the three-dimensional network is perfect (a crystal), then there is no way any chemical or physical disruption can affect the network.

By contrast, sp2 carbon atoms have extra electron density between each two carbons. Any strong chemical agent or physical intervention can use that electron density to break in and distrupt the structure. So graphite and nanotubes arenot stronger than diamond.

After all, we write with graphite pencils and leave black marks all over the page. How stable is that?

Answer 3

If bond length is any indicator, the individual sp2 bonds in graphite *are* stronger than the sp3 bonds in diamond.
http://www.citycollegiate.com/carbonsiliconIXa.htm
The difference in hardness of the bulk material is in the three-dimensional arrangement of the bonds. Graphite is organised into sheets; each sheet is not really bonded to the one below, so adjacent sheets slide against each other. If you "rip" graphite along one of these planes, there's not much resistance. If you "rip" perpendicular to it, it will be tougher, but still the layers will deform and slide.
In diamond, there are no such layers. Every which way you try to "rip" it, you have bonds going up-down-left- right-back-and-forth. Each atom is anchored to four neighbors (not just the 3 in graphite), so trying to move even one is difficult -- you're trying to drag so many others with it, in every direction (not just in the same plane or layer as in graphite).

Prof. Geim in Manchester, England has been "peeling" off single layers of graphite ("graphene") using Scotch tape. These sheets are just one atom thick, so you can just imagine how tough each sheet must be, if you can handle it with something as clumsy as sticky tape.
http://www.nytimes.com/2007/04/10/science/10grap.html?ex=1333857600en=733784a971ea5f29ei=5088partner=rssnytemc=rss

Answer 4

Graphite and diamonds (and also coal) are made from carbon, and diamonds have the strongest bonds in the world. Graphite has strong bonds too but not in all directions... meaning it flakes off when rubbed in a certian direction, which is why pencil "leads" are made from this.