Are Diamonds Forever? A Chemical Investigation

A diamond is forever…at least that’s how the advertising slogan goes. Many chemists know this saying is not entirely true, because diamonds are converted to graphite under normal conditions:

C(s, diamond) à C(s, graphite)                     DG^o = -2.9 kJ mol^-1                 Equation 1

However, the conversion process is extremely slow because an enormous activation energy barrier exists for this process: about 370 kJ mol^-1.¹ Thus, the conversion occurs extremely slowly – over billions of years – which allows us to enjoy the beautiful sparkle of diamonds in earrings, necklaces, and engagement rings.

Nevertheless, the fact that diamonds can spontaneously convert to graphite can be seen by inspecting the phase diagram for carbon,² which clearly shows that graphite is the stable phase of carbon at atmospheric pressure and ambient temperature. In addition, thermodynamic values (Table 1) can be used to show that diamond spontaneously converts into graphite: DG for this process is negative, and the process is favored both enthalpically (DH^o = -1.8 kJ mol^-1) and entropically (DS^o = +3.2 J mol^-1 K^-1).

Table 1 - Thermodynamic values for substances of interest in this work.

However, many chemists are not aware that diamonds can be converted to carbon dioxide by simply burning them.^2,3 That’s right: diamonds are not forever because they can be burned:

C_{(s, diamond)} + O2_(g) à CO_2 (g)                     DG^o = -397.3 kJ mol^-1                 Equation 2

Methods of burning diamonds that are simple enough to conduct in the classroom exist.² For example, check out the following video:

The combustion of diamonds (Equation 2) allows for an easier “destruction” of diamonds than conversion to graphite (Equation 1) for a number of reasons. First, the activation energy for the combustion of diamond is only in the neighborhood of 150 – 250 kJ mol^-1:¹ much lower than the activation energy for conversion to graphite. Second, combusting diamonds can be accomplished at temperatures of 600^oC to 900^oC, which are easily obtained by heating with a Bunsen burner or blow torch. Finally, the combustion of diamond is much more thermodynamically favored than conversion of to graphite, with DG^o = -397.3 kJ mol^-1 for the former and DG^o = -2.9 kJ mol^-1 for the latter. Both entropy (DS^o = +6.2 J mol^-1 K^-1) and enthalpy (DH^o = -395.3 kJ mol^-1) favor the combustion reaction, which is quite exothermic (Table 1).

Diamonds are not forever – and now you have some chemistry demonstrations that you can perform in your classroom to prove it!

References:

1. Chen Y., Theoretical Study of Material Removal Mechanism in Polishing of Polycrystalline Diamond Composites, Ph. D. Thesis, University of Sydney, 2007. 

2. Bundy, F.P., Bassett, W.A., Weathers, M.S., Hemley, R.J., Moa, H.K., Goncharov, A.F., The Pressure-Temperature Phase Transformation Diagram for Carbon; Updated through 1994, Carbon, Vol 34, No. 2, pp 141-153. 1996.

3. Miyauchi, T. and Kamata, M. Classroom Demonstration: Combustion of Diamond to Carbon Dioxide Followed by Reduction to Graphite, Journal of Chemical Education, 2012, 89 (8), 1050-1052.