Why does diamond have a higher boiling point than graphite?

Question

answer with reference to structure and bonding

a little more detail would be helpful...answer explaining WHY the BP is greater than graphite's.
thnx

Answer 1

The physical structure of a diamond and graphite are different, even though they are both pure carbon.

Graphite is a stack of carbon rings which are not attached to each other. That's why it slides easily and is a good lubricant and used in pencils.

Diamonds are a stack of carbon rings with cross members that are all connected. Its called covalent network bonding. This makes them very hard.

For the same reason, it will be harder to pull diamond molecules apart compared to graphite, and so the BP will be higher.

Answer 2

Boiling Point Of Diamond

Answer 3

A diamond is a transparent crystal of tetrahedrally bonded carbon atoms and crystallizes into the face centered cubic diamond lattice structure. Diamonds have been adapted for many uses because of the material's exceptional physical characteristics. Most notable are its extreme hardness, its high dispersion index, and extremely high thermal conductivity (900 – 2320 W/m K), with a melting point of 3820 K (3547 °C / 6420 °F) and a boiling point of 5100 K (4827 °C / 8720 °F).[5] Naturally occurring diamond has a density ranging from 3.15 to 3.53 g/cm³, with very pure diamond typically extremely close to 3.52 g/cm³.

In graphite each carbon atom is covalently bonded to three other surrounding carbon atoms. The flat sheets of carbon atoms are bonded into hexagonal structures. These exist in layers, which are not covalently connected to the surrounding layers. Instead, different layers are connected together by weak forces called van der Waals forces much like those of mica.

The unit cell dimensions are a = b = 2.456 ångströms = 245.6 picometers, c = 6.694 Å = 669.4 pm. The carbon-carbon bond length in the bulk form is 1.418 Å (141.8 pm), and the interlayer spacing is c/2 = 3.347 Å (334.7 pm).


Crystal structure of graphiteEach carbon atom possesses an sp² orbital hybridisation. The pi orbital electrons delocalized across the hexagonal atomic sheets of carbon contribute to graphite's conductivity. In an oriented piece of graphite, conductivity parallel to these sheets is greater than that perpendicular to these sheets.

The bond between the atoms within a layer is stronger than the bond of diamond, but the force between two layers of graphite is weak. Therefore, layers of it can slip over each other making it soft.

Answer 4

Diamond has a internal sp3 bonding structure where each carbon bonds to 4 others. The diamond exists as an extended tetragonal array.

Graphite has an internal sp2 bonding structure, where each carbon bonds to 3 other carbons. Moreover, the bonds cause the extended graphite structure to be in unilayer sheets with only occasional bonding across sheets.

Answer 5

Diamond and graphite are allotropes of carbon. They differ in their physical and chemical properties because of differences in the arrangement and bonding of the carbon atoms. Diamond is denser than graphite (3.51 vs. 2.22 g per cubic centimetre), but graphite is the more stable, at 300 K and 1-atm pressure. From the densities it follows that to transform graphite into diamond, pressure must be applied, and from the thermodynamic properties of the two allotropes it can be estimated that they would be in equilibrium at 300 K under a pressure of about 15,000 atm. Of course, equilibrium is attained extremely slowly at this temperature. At high temperature in vacuum, diamond transforms into graphite, which sublimes at 1500 degrees C.

Thermodynamically, it is said that two states are in equilibrium if they have the same free energy. Let us examine this criterion for a reaction such as


C (diamond) = C (graphite)

The difference in free energies of formation, DDG, at 300 K and 1 atm is


DG (graphite) - DG (diamond) = -2.9 kJ per mol

Since DDG is negative, the reaction proceeds as written, and we can conclude that graphite is thermodynamically the more stable form.

Interestingly this also accounts for the conversion of graphites into diamonds and vice versa.Thus,Carbon is polymorphic, meaning it can have more than one form. Carbon atoms can bond together to form either graphite or diamonds, or something else, depending on the organization of its atomic structure. At all temperatures, graphite is the more stable state for carbon. Under extremely high temperatures and pressures the graphite can be converted to diamonds. The bonds between the carbon atoms are shorter in diamonds than the long bonds in graphite. The extreme pressure causes the compression, or shortening or the bonds, and then the rearrangement of atoms necessary to convert graphite to diamonds. This creates a more efficient packing of the atoms. It is true that, according to the laws of thermodynamics, a diamond spontaneously changes to graphite under atmospheric pressure. The reason that diamonds can be found on the Earth's surface is because they are metastable. The change from diamonds to graphite takes place so slowly that it cannot be detected.

Answer 6

There are more covalent bonds to break in diamond - 4 as opposed to about 3.5 per C atom.
You can then refer to the correct structures of diamond and graphite, of course.

Answer 7

use this website to help you:

(and why would you boil diamonds anyway??