Chemistry - Thermochemistry Question with Solution | TestHub

The reaction given is:

$$C(graphite)\to C(diamond)\Delta H=+1.9kJ$$

This means that converting graphite to diamond requires 1.9 kJ of energy per mole of carbon. In other words, diamond has 1.9 kJ/mol MORE energy than graphite. This also means graphite is more stable than diamond.

When both diamond and graphite are burned, they form CO₂:

$$C(diamond)+O_2(g)\to C{O}_2(g)\Delta H_1$$

$$C(graphite)+O_2(g)\to C{O}_2(g)\Delta H_2$$

Since diamond has a higher energy content than graphite, when diamond is burned, it will release LESS heat (in magnitude, considering the negative sign of enthalpy change) than when graphite is burned.

The difference in heat released will be equal to the difference in their enthalpies of formation, which is 1.9 kJ/mol.

We are given 6 g of each allotrope. The molar mass of carbon is 12 g/mol.

Therefore, the number of moles of carbon is:

$$n=\frac{6g}{12g/mol}=0.5mol$$

The difference in heat released when 0.5 mol of diamond is burned compared to 0.5 mol of graphite is:

$$\Delta H=0.5mol\times 1.9kJ/mol=0.95kJ$$

Since diamond has higher energy, burning diamond will liberate LESS heat than burning graphite. Therefore, the heat liberated in the first case (burning diamond) is less than in the second case (burning graphite) by 0.95 kJ.