Graphite transforms into diamond

Graphite transforms into diamond upon the application of pressure P (in atm) and temperature T (K). This relationship is expressed by the following equation ... 

Like graphite, C60 can be transformed into diamond, but the process requires less stringent conditions. It has also been found that Cso becomes a superconductor at low temperature. Another interesting characteristic of Cso is that when it is prepared in the presence of certain metals, the Cso cage can enclose a metal atom. In some cases, other materials can be enclosed within the C60 cage in a "shrink wrapped" manner to form "complexes" that are described as endohedral. It has also been possible to prepare metal complexes of Cso that contain metal-carbon bonds. A compound of this type is (C6H5P)2PtC60. 

At 30 C and 1 bar, graphite is about 2900 J/mol more stable than diamond. Both forms are in equilibrium at 300° C and 15000 bar. At 2700°C and pressures over 125000 bar, graphite can be transformed into diamond. However, the reaction is very slow and requires acceleration by catalysts (Cr, Fe, Pt). 

At room temperature graphite is more stable than diamond however, at high pressures and temperatures, graphite can be transformed into diamond if molten chromium, iron, or nickel is present as a catalyst. Although these synthetic diamonds are not gem quality, they are useful in drills and saw blades. 

Fullerene-Diamond Transformation. The rapid compression of Cqo powder, to more than 150 atm in less than a second, caused a collapse of the fullerenes and the formation of a shining and transparent material which was identified as a polycrystalline diamond in an amorphous carbon matrix.O Thus the fullerenes are the first known phase of carbon that transforms into diamond at room temperature. Graphite also transforms into diamond but only at high temperatures and pressures (see Ch. 12, Sec. 3.0). 

Elucidation of the phase relationships between the different forms of carbon is a difficult field of study because of the very high temperatures and pressures that must be applied. However, the subject is one of great technical importance because of the need to understand methods for transforming graphite and disordered forms of carbon into diamond. The diagram has been revised and reviewed at regular intervals [59-61] and a simplified form of the most recent diagram for carbon [62] is in Fig. 5. 

Diamond is transformed into graphite when heated by a powerful electric current between carbon poles, and both diamond and graphite can be indirectly converted into charcoal. The artificial production of the diamond, however, is a more difiticult process but the late Professor Moissan succeeded in effecting it, so far as very small diamonds are... 

It has been reported that when Ceo is rapidly and nonhydrostatically compressed above 20 GPa at room temperature, it transforms into polycrystalline diamond [524]. Although Ceo can be considered as a folded graphite sheet, we must take into account that in the pentagons there is an important tetrahedral distortion making the transformation of Ceo into diamond likely easier than the HP-HT conversion from graphite, and it is possible to use this reaction for industrial production of diamonds. 

The second type is simple phase transitions in which one phase transforms into another of identical composition, e.g., diamond graphite, quartz coe-site, and water ice. This type sounds simple, but it involves most steps of heterogeneous reactions, including nucleation, interface reaction, and coarsening. 

Carbon is likely to congeal to high-molecular-weight polymers as H2 distills off. In extraterrestrial environments, we expect lower hydrocarbons eventually to transform into pure carbon, either diamond (in which all the carbons are singly bonded to other carbons), fullerenes and graphite (in which each interaction between a pair of carbons is the approximate equivalent of 1.5 bonds), or carbon bonded to other elements that cannot be converted to a volatile form. 

Diamond, graphite, and the fullerenes differ in their physical and chemical properties because of differences in the arrangement and bonding of the carbon atoms. Diamond is the densest (3.51 vs 2.22 and 1.72 g cm-3 for graphite and Cw, respectively), but graphite is more stable than diamond, by 2.9 kJ mol-1 at 300 K and 1 atm pressure it is considerably more stable than the fullerenes (see later). 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 —15,000 atm. Of course, equilibrium is attained extremely slowly at this temperature, and this property allows the diamond structure to persist under ordinary conditions. 

Although graphite is thermodynamically more stable than diamond at 25°C and 1 atmosphere, a diamond will not transform into graphite, even over a period of thousands of years. Which of the following correctly explains this observation ... 

Although the enthalpy of formation under standard conditions for diamond is l.OkJmoT higher than that for graphite, the first does not spontaneously transform into the latter, but rather represents a metastable modification under normal conditions. In contrast to graphite it is not at all anisotropic regarding any property. 

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