Wait, why is graphite more stable than diamond?

Astudentinneed

that doesnt make sense

Reply 1

gooner88

no clue... :confused:

Reply 2

Rocky27

despite both being macromolecules, graphite has van der waals forces in between adjacent layers in its structure (in addition to the many covalent bonds), while diamond doesnt

Reply 3

dagorian

but i could have sworn the van derwaals forces in between the planes are weak and surely many covalent bonds of the diamond structure more than makes up for it

Reply 4

Rocky27

tru the vdw forces between layers are weak, but in such a big molecule they make a big difference as the strength of them increases with size....i hope, im not 100% sure but it seems logical

Reply 5

ma3eeni

Rocky27

tru the vdw forces between layers are weak, but in such a big molecule they make a big difference as the strength of them increases with size....i hope, im not 100% sure but it seems logical

noo ..
the melting point of graphite is larger than diamond ...
the melting point here is the energy required to break the carbon atoms WITHIN the layers ...
vander vaals forces are still weak.

Reply 6

gyrase

My attempted explanation;
Diamond does not contain any delocalised electrons. Graphite containes one delocalised electron per carbon. These cause greater attraction between carbon atoms hence giving stronger bonds, more stability to the structure.

Reply 7

mysteryman

Graphite has greater van der Walls forces. Because of the oscillating delocalised electrons which induce temporary dipoles in graphite increasing attraction and stability. Diamond all delocalised electons are used in covalent bonding.

Reply 8

Orizon

Isn't the answer supposed to be related to how diamond is at a lower energy level with respect carbon in its standard state compared to graphite with respect to carbon in its standard state, and nothing to do with bonding at all?

Reply 9

sea tea

seriously, wikipedia is your friend:

Thermodynamic stability of Diamond
At surface air pressure (one atmosphere), diamonds are not as stable as graphite, and so the decay of diamond is thermodynamically favorable (ΔG = −2.99 kJ / mol). Diamonds will burn at approximately 800 degrees Celsius, providing that enough oxygen is available. This was shown in the late 18th century, and previously described during Roman times. However, owing to a very large kinetic energy barrier, diamonds are metastable; under normal conditions, it would take an extremely long time (possibly more than the age of the Universe) for diamond to decay into graphite.