# Why don't decomposition reactions form pure elements as products?

For example, if I had a reaction like this:

$$\ce{NaHCO3 -> Na2CO3 + H2O + CO2}$$

• Why does it not break down all the way down to its elements?
• What makes it form such "intermediate" products, stopping at water, carbon dioxide, etc.?
• If you heat sodium bicarbonate strongly you can convert it to $\ce{Na2O}$. In general the reaction favors the most thermodynamically stable chemical.
– MaxW
Feb 21 '17 at 1:28

Using your example, you simply have not given the system enough energy to do so. Water and $\ce{CO2}$ are very stable compounds under ordinary conditions of temperature and pressure, and thus are the preferred products under standard conditions for that reaction.
If you were to heat the system further, further decomposition would occur. For example, at higher temperatures you will convert $\ce{Na2CO3}$ into $\ce{NaO}$ and more $\ce{CO2}$. Under extreme conditions, i.e. thousands of degrees K, you will ultimately have enough energy to break all of the chemical bonds and then you will be left with the elements.