# Why doesn't sodium carbonate decompose?

I was asked to find the hypothetical $\Delta H$ and $\Delta S$ of the decomposition of sodium carbonate and then asked to explain why it doesn't decompose.

I calculated the enthalpy and entropy change and from that I calculated the Gibbs free energy change. I was hoping to get a positive answer, but actually I got a negative answer, indicating that the process is actually spontaneous.

So is the reason that sodium carbonate doesn't decompose due to its kinetic properties, i.e. the reaction is too slow?

• What is the decomposition pathway? Don't leave us in the dark! But in general if it's thermodynamically favoured but doesn't happen then it's a kinetic factor. – orthocresol Feb 20 '16 at 20:33
• Your $\Delta G$ must be wrong. Sodium carbonate is thermodynamically stable. – Ivan Neretin Feb 20 '16 at 20:36
• I calculate a positive standard free energy change for the decomposition reaction. – Chet Miller Feb 20 '16 at 21:31
• Okay, thanks for all the help. I double checked my values I still get a negative value. The question probably gave the wrong values for the enthalpy and entropy change. – Nanoputian Feb 20 '16 at 22:56
• OK. How about telling us what you got for the free energies of formation of each of the products and the reactant? – Chet Miller Feb 21 '16 at 0:13

From the values on the Wikipedia page for sodium carbonate, the Gibbs free energy of formation is negative and thus, formation is thermodynamically preferred over decomposition at $\pu{298K}$. Your given values seem to be wrong after all.
By the way, $\Delta H$ and $\Delta S$ should be the negative of those given on that page for your case, making $\Delta G$ a positive number.