What is the enthalpy change of graphite to diamond?

Selected ATcT enthalpy of formation based on version 1.122 of the Thermochemical Network

Enthalpy of combustion of carbon graphite and carbon diamond is -293.5(kJ/mol) And – 670(kJ/mol) respectively.

The enthalpies of combustion of graphite and diamond are 393.5 kJ And 395.4 kJ respectively.

Color

Solution : Graphite is most stable of carbon, hence its energy is lower than that of diamond and entropy of graphite is also lower than that of diamond.

The enthalpy of both graphite and diamond is taken to be Zero, being elementry substances.

Enthalpy is the heat content of a system. The enthalpy change of a reaction is Roughly equivalent to the amount of energy lost or gained during the reaction. A reaction is favored if the enthalpy of the system decreases over the reaction.

Answer: The enthalpy change in the conversion of 10 g graphite in diamond is 0.375 kcal. The enthalpy change in the conversion of 10 g graphite in diamond is 0.375 kcal.

I’m not sure the numbers are right in this quiz I found, but if you run the numbers on the entropy at different temperatures, the entropy does change sign, i.e. At phase transition temperatures, the entropy of diamond is higher than that of graphite (for some reason).

According to thermodynamics, the reaction from diamond to graphite is spontaneous and favorable. However, because kinetics rather than thermodynamics is controlling this reaction, it occurs extremely slowly. So, diamond is kinetically stable, but thermodynamically unstable.

From the phase diagram of carbon, it can be seen that diamond is the thermodynamically favored allotrope under geological conditions of high pressure, but at ambient conditions, graphite is the more stable allotrope, and Diamond spontaneously converts to graphite.

Solution : `DeltaH = DeltaU` during a process which is carried out in a closed vessel `( DeltaU = 0 )` or number of moles of gaseous products `=` number of moles of gaseous reactants or the reaction does not involve any gaseous or product.

The conversion of diamond into graphite is an endothermic reaction.

Greater entropy of graphite is related to its structure as graphite is less compact and rigid than diamond. ΔHf∘ for graphite is zero, but the ΔHf∘ for diamond is 2kJ/mol. That is because Graphite is the standard state for carbon, not diamond.

The enthalpy difference between graphite and diamond is too large For both to have a standard enthalpy of formation of zero. To determine which form is zero, the more stable form of carbon is chosen. This is also the form with the lowest enthalpy, so graphite has a standard enthalpy of formation equal to zero.

A pure element in its standard state Has a standard enthalpy of formation of zero.

Specific enthalpy

It can be expressed in other specific quantities by H = u + pv, where u is the specific internal energy, p is the pressure, and v is specific volume, which is equal to 1/ρ, where ρ is the density.

The conversion of diamond into graphite is an Endothermic reaction.

1 The enthalpy change of atomisation of diamond is smaller than that of graphite. 2 The bond energy of the C–C bonds in graphite is greater than that in diamond. 3 The enthalpy change of combustion of diamond is greater than that of graphite.

According to thermodynamics, the reaction from diamond to graphite is spontaneous and favorable. However, because kinetics rather than thermodynamics is controlling this reaction, it occurs extremely slowly. So, diamond is kinetically stable, but thermodynamically unstable.

The conversion of carbon nanofibers and nanotubes into diamond nanofibers involves Melting in a super undercooled state using nanosecond laser pulses, and quenching rapidly to convert into phase-pure diamond. The conversion process occurs at ambient temperature and pressure, and can be carried out in air.